

ITS USES 




»N AND 



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mmm^m 

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nass QTi 33 
Book • M £ 3 
GopigM? 

COPYRIGHT DEPOSIT. 




Lavoisier (1743-1794) 

The great French chemist, Antoine Laurent Lavoisier, the founder 
of modern chemistry, is here shown in his laboratory conducting 
the famous experiments by which he proved the true nature of 
combustion. He was guillotined during the French Revolution 
because of bis connection with the government 



CHEMISTRY AND ITS USES 



A TEXTBOOK FOR SECONDARY SCHOOLS 



BY 

william Mcpherson 

AND 

WILLIAM EDWARDS HENDERSON 




GINN AND COMPANY 

LOSTON • NEW YORK • CHICAGO • LONDON 
ATLANTA • DALLAS • COLUMBUS • SAN FRANCISCO 



COPYRIGHT, 1922, BY WILLIAM MCPHERSON AND 

WILLIAM E. HENDERSON 

ENTERED AT STATIONERS' MALL 

ALL RIGHTS RESERVED „- 

SO 



322.5 



Q^V 3 



TEfte gtftengum jgregg 

GINN AND COMPANY • PRO- 
PRIETORS • BOSTON • U.S.A. 



ICI.A674 5H0 



PREFACE 

Every teacher of chemistry in the high school realizes that 
the watchword of the present day is the practical rather than 
the theoretical, the application rather than the abstract princi- 
ple, the pictorial rather than the descriptive. Just how far this 
tendency should be followed each author and each teacher 
must decide for himself ; this volume represents the opinion 
of the present authors. The text abounds in the practical 
applications of chemistry in the arts and industries as well 
as in everyday life, and no effort has been spared to have the 
illustrations as attractive, as instructive, and as accurate as 
possible. The authors have also kept before them the prac- 
tical aim of presenting the occupation of the chemist as one 
that is very attractive to a boy who is thinking about what 
he will do in the world. 

To the teachers using this book the authors wish to empha- 
size the conviction that the practical applications of chemistry 
have a place in high-school instruction largely in as far as they 
are used to illustrate the principles of the science and the way 
in which pure chemical knowledge can be turned to the uses 
of society. The main object of the course in chemistry must 
always be to train young people to think and to imagine in the 
realm of chemical facts and laws, and the teacher who finds 
at the end of the course that his pupil has acquired what 
seems to be a fund of useful information, but has little ability 
to think for himself how he would solve a simple chemical 
problem, should feel dissatisfied with his effort. 

The utilization of chemical knowledge in the problems that 
developed during the World War is a fascinating story, and 



vi CHEMISTRY AND ITS USES 

many paragraphs have been devoted to this theme. It is hoped 
that the student will be led to see how chemistry can solve 
new problems when new conditions arise and will not rest 
content in merely contemplating the chemical triumphs of 
the war, great as these were. 

Throughout the text a great many items of interesting in- 
formation as well as supplementary explanation have been 
thrown into subordinate type. It is believed that the matter 
in large type is complete in itself and nowhere dependent on 
the subordinate paragraphs. However, it is thought that so 
much of interest will be found in these paragraphs that the 
teacher can easily induce the student to read them even if 
they are regarded as excess material. Directions for laboratory 
work will be found in a separate volume. 

The authors are indebted to a large number of firms for 
assistance in securing photographs. Among these are the 
following: the Corning Glass Works, the American Rolling 
Mill Company, the Eastman Kodak Company, the Anaconda 
Copper Mining Company, the American Cyanamide Company, 
the H. V. Bretney Leather Company, the Goodrich Rubber 
Company, and the Koppers Company. Many high-school 
teachers have offered valued suggestions, and the authors 
desire to express their hearty appreciation to all these. We 
are especially indebted to our colleague, Mr. William Lloyd 
Evans, to Mr. George M. Strong, and to Mr. Robert W.Collins, 
instructors in chemistry in the East High School of Columbus, 
Ohio, and to Mr. Charles S. Pease and Mr. Julius Stone, Jr., 
of the Department of Chemistry, the Ohio State University. 
Mr. Strong and Mr. Stone not only read the proof but offered 
many suggestions as well. mTTT , , „m„^„ 

J && THE AUTHORS 



CONTENTS 

CHAPTER PAGE 

I. Chemistry and the Work of the Chemist . . 1 

II. Matter and its Classification 6 

III. Oxygen 14 

IV. Hydrogen 30 

V. How Gases Act: how they are Made Up . . 11 

VI. Matter and Energy 50 

VII. Compounds of Hydrogen and Oxygen: Water and 

Hydrogen Peroxide 61 

VIII. Nitrogen and the Rare Elements Argon, Helium, 

Neon, Krypton, Xenon 76 

IX. Molecular Weights; Atomic Weights .... 82 

X. Formulas; Equations; Solution of Problems . 91 

XI. Carbon and its Oxides 102 

XII. Valence 118 

XIII. The Air 122 

XIV. Solutions 130 

XV. Chlorine; Hydrogen Chloride; Hydrochloric 

Acid; Acids and Salts 13G 

XVI. Sodium; Sodium Hydroxide; Bases 150 

XVII. Ionization 156 

XVIII. Compounds of Nitrogen 163 

XIX. Reversible Reactions; Equilibrium 170 

XX. Sulfur and its Compounds 181 

XXI. The Periodic Law 107 

XXII. The Chlorine Family 203 

XXIII. Hydrocarbons: Petroleum . 211 

XXIV. Fuels; Electric Furnaces; Flames 219 

XXV. Coal-Tar Compounds 282 

XXVI. Carbohydrates and Textiles 238 

XXVII. Alcohols; Preseryatiyes 252 

XXVIII. Organic Acids and their Derivatives ; Proteins 259 

XXIX. Foods 266 

XXX. The Phosphorus Family 275 

vii 



viii CHEMISTRY AND ITS USES 

CHAPTER PAGE 

XXXI. Silicon and Boron 285 

XXXII. Colloids — the Chemistry of Very Small Par- 
ticles 293 

XXXIII. The Metals 301 

XXXIV. The Sodium Family 305 

XXXV. Soap ; Glycerin ; Explosives 318 

XXXVI. The Calcium Family 327 

XXXVII. Soils and Fertilizers 338 

XXXVIII. The Magnesium Family 344 

XXXIX. Aluminium 353 

XL. Silicates and their Commercial Applications 366 

XLI. The Iron Family 372 

XLII. Copper, Mercury, and Silver 390 

XLIII. Tin and Lead 401 

XLIV. Manganese and Chromium 411 

XLV. Platinum and Gold 416 

XLVI. The Story of Radium 421 

XLVII. Some Applications of Rarer Elements . . . 428 

APPENDIX 

Chemical Library . 431 

Thermometers 434 

Densities and Melting Points of Some Common Ele- 
ments 436 

Table of Solubility of Various Solids 436 

Tension of Aqueous Vapor expressed in Millimeters 

of Mercury 436 

INDEX 437 

REFERENCE TABLES 

List of the Elements, their Symbols and their 

Atomic Weights Facing inside of back cover 

Weight in Grams of 1 Liter of Various Gases under 
Standard Conditions and Boiling Points under 

Pressure of 760 Millimeters . . Inside of back cover 

Displacement (Electrochemical) Series Inside of back cover 
Relation between English and Metric Constants 

Inside of back cover 



CHEMISTRY AND ITS USES 

CHAPTER I 
CHEMISTRY AND THE WORK OF THE CHEMIST 

The general field of chemistry. At the beginning of any new 
study it is well to get at least a general idea of the work ahead. 
So we may say at the outset of the course in chemistry that 
we shall be interested in the changes that the material things 
around us undergo. In some cases we notice that things change 
because they wear out or are broken, but they undergo no 
change in composition. Thus a dish may break and be valueless 
as a dish, but the pieces will have the same composition as the 
original dish. We shall have but little interest in such changes. 

In most of the changes taking place around us it is evident 
that the very nature of the matter is changed. Thus, nearly 
all the metals tend to rust or corrode ; coal and wood burn 
to form ashes and invisible gases ; the constituents of the soil 
and the air are built up into a vast variety of living organ- 
isms ; the food we eat is changed into fat and bone and muscle. 
In all these examples the materials undergo a change in 
composition, and it is in such changes that the chemist is 
especially interested. The science of chemistry deals with all 
those changes that result in altering the composition of materials. 

Relation of chemistry to the other sciences. Since all the 
studies we call natural sciences, such as physics, biology, 
geology, and physical geography, in part deal with changes 
in the composition of materials, it is evident that chemistry 
is fundamental to all of them. We cannot study electrical 

l 



2 CHEMISTRY AND ITS USES 

changes, or digestion, or rock disintegration without a knowl- 
edge of the exact changes that take place as well as of the 
general result of the changes. Physics and chemistry are 
really one science. Physics deals chiefly with what we call 
energy, — that is, with motion, heat, light, sound, electricity, — 
while chemistry deals with changes in the composition of matter. 
But almost every occurrence in nature that causes a change in 
matter causes a change in energy as well, and we cannot sepa- 
rate the two from each other. Moreover, the chemist is often 
interested quite as much in the energy set free when matter 
undergoes a change as in the change itself. For example, he 
is often called upon to measure the heat given off by burning 
a given sample of coal, for it is this that determines the value 
of the coal, and not the products that are formed. 

The alchemists. In the earliest days of chemistry the chief 
chemical occupations were those of producing the metals from 
their ores, making glasses and enamels, and dyeing fabrics. 
There was no understanding of what happens in these processes, 
and the only guide was experience. In time the idea took a 
strong hold on the minds of these workers that the various 
metals are not really different things, but are merely stages in 
the purification of the one metal gold, and they thought that it 
ought to be possible to change all metals into gold in stages. 
These early chemists were called alchemists. It took centuries 
of work for them to become convinced that one metal cannot 
be changed into another. But in their efforts the alchemists 
found out many new facts, made many new compounds, and 
devised new processes, all of which had great practical value. 

Almost at the time when the American Revolution for inde- 
pendence was starting, the great discovery of oxygen was 
made (1774), and the nature of the process we call burning or 
combustion became understood. This provided a sound foun- 
dation for the understanding of chemical processes of all kinds, 
and chemistry developed rapidly into a true science. 



CHEMISTRY AND THE CHEMIST 3 

Nature supplies raw materials only. Very few of the mate- 
rials found in nature are suited to the needs of man advanced 
beyond the most primitive stages of development. From the 
beginning he has had to match his wits against nature to gain 
materials better suited to his needs. At first he relied upon 
merely sorting out the best material and rejecting that of poor 
quality. For example, he pulled out the fine fibers from the 
stem of the flax plant, and he hammered or melted the small 
particles of gold found in sands into larger pieces. But he 
could not go very far in this way alone. 

New materials from plants and animals. Man's most pressing 
need has always been for food and clothing. When natural 
species failed to supply his needs, he set about to improve on 
nature. The grains and vegetables and fruits we now grow for 
food and the plants that give us textile fibers bear little re- 
semblance to the original plants of nature. The original beet 
had just enough sugar in it to suggest an idea ; the modern 
beet has as much as 16 per cent. The cotton plant of 
nature would not be recognized beside the long-fibered plant 
of Arizona today. So, too, with the animals that supply us 
with meat, fat, milk, and wool. They have been very highly 
developed from primitive stock. 

Very little of this work has been done by the chemist. Yet 
it must not be overlooked that each plant and animal is doing 
just what the chemist does in his laboratory. It takes the raw 
materials supplied by nature and builds them up by chemical 
changes into very diversified products. All living organisms 
are therefore chemical laboratories — much more wonderful 
and efficient than the ones we build. So in improving a plant 
or animal we are really improving a chemical process. 

New materials from inanimate products. The chemist for the 
most part works with matter that is not living ; namely, with the 
minerals supplied by nature and with the products of life such 
as fats, oils, sugars, starch, and hundreds of similar products. 



4 CHEMISTRY AND ITS USES 

From these he has developed countless new materials for 
human needs. From the rocks he has prepared a large array 
of useful metals and such indispensable things as lime and 
cement. From coal he has manufactured thousands of dyes, 
medicinal preparations, fertilizers, paints, and other useful but 
less familiar materials. From petroleum and wood and the 
saps of trees, from milk and hides and hoofs, he has fashioned 
the materials that make the physical comforts surrounding the 
lives of even the poorest luxurious when compared with the 
possessions of the kings of the past. 

Chemical industries. All the industries that transform a 
raw material of nature into a finished product for human use 
are essentially chemical industries. They must rest on the prin- 
ciples of chemical science, and they must be supervised by 
someone who understands chemistry. It is only as our knowl- 
edge of chemistry increases that these industries can be con- 
ducted more economically and expanded into more efficient 
and diversified ones. The industrial advancement of a country 
can be judged fairly well by the extent to which chemistry is 
cultivated in its schools and universities. 

The work of the chemist. Some chemists will always be 
chiefly interested in adding to our knowledge of chemistry as 
a science. Others will have as their greatest interest some 
definite industrial plant (Fig. 4). Such a one must supervise 
the process used in the industry and devise improvements. He 
must find out the composition of each new lot of raw material 
and modify the process accordingly. He must be able to guar- 
antee the composition and the properties of the output of the 
factory so that its value may be known. He must study the 
possibilities of making something useful out of the waste prod- 
ucts of the business. For example, he has taken the formerly 
useless cotton seed and from it has made most valuable oils, 
fats, soaps, and cattle feeds. Since industry has such a wide 
range of raw materials employed, there is almost no limit to 



Fig. 1. Ira Remsen 






■■PP^H 




(1846- ) 


J- 




For many years director of 


'*% IBH 




chemical research at the Johns 


i 




Hopkins University, later presi- 


Ijl 




dent of that institution ; known 


\ • ^-'"-"v ■ 




equally well as a great chemist, 


\ 




an inspiring teacher, and an edu- 


Jm ^H " PV JH 




cational leader; editor of an 


.^^^^k ^1 St • ^9 




important chemical journal and 


^M ^k ^wl i^H 




the author of an important series 


Ik m Ml '^H 




of books; president of the 


K « A 




American Chemical Society 


mL^* T tji ' *'\ 




(1902) ; member and officer in 


^^^r V ^^R^fc& » ■ * 




many of the important scientific 


wt H| 1 




societies of this country as well 


■fl^^. Hlf M i 




as foreign member of many 


^EVfek^^H«H ^m 




European societies 


BBW1 ■ 




Fig. 2. Edgar Fahs Smith 






(1856- ) 






For many years director of 






chemical research at the Uni- 






versity of Pennsylvania and later 






president of that institution; 






noted for his contributions to 


' $t ■ 




electrochemistry and to our 


R\*^Ii3Hft^ 




knowledge of the rarer ele- 


h^BhBl 




ments; a writer of fascinating- 


»^*HBb ife^ 




interest on the early history of 




chemistry in the United States: 


■ JpKifW^/;/' ; 1 ;- 




president of the American Chemi- 


wIMmm Jm 




cal Society (1895, 1921, and 1922) ; 


: J? ~'*-^r " ■■• '*'' ; X'-^'«' ; 




member and officer in many sci- 


Hlhfl^^l 




entific societies in this country 


'...'■■"}■ :M* ■•'/-."*. ' ,' - 




and abroad 


aHHHHi 












Fig. 3. An alchemist in his laboratory 




Fig. 4. A typical laboratory of a modern automobile plant 



CHEMISTRY AXD THE CHEMIST 5 

the variety of the work of the chemist. The chemist also 
helps to preserve the health of the people by finding out the 
kinds and amounts of food best adapted to our bodies, by 
devising methods for purifying our water supplies and for 
disposing of sewage, and by preparing various substances that 
are useful in combating diseases. 

Chemical knowledge incomplete. While we seem to know a 
good deal about changes in matter, we have really just made 
a good beginning. For example, our knowledge of what takes 
place in growing plants and animals is very limited, though 
we are learning fast. Plants manufacture sugar and starch 
and cellulose chiefly from air and water, but we cannot do 
this in the most expensively equipped laboratories, nor do 
we understand how the commonest plant accomplishes this 
seeming miracle. Even in well-understood processes there is 
usually some detail that is not yet clear. 

The future of chemistry. There is a wealth of unsolved 
problems of vast importance to human comfort and happiness 
to attract the student to the life pursuit of chemistry. For 
example, we are rapidly using up our reserves of easily 
available raw materials, such as coal, oil, gas, wood, and 
metal ores, and substitutes must be found. We are exhaust- 
ing our fertile lands, and we must learn better how to keep 
them in productive condition. Every line of material progress 
along which we are moving is full of these problems, and as 
a result the number of those engaged in chemical occupa- 
tions is increasing very fast. In every progressive country 
the national chemical society is one of the largest of the 
scientific organizations. The American Chemical Society is 
the largest scientific society in the world, having more than 
15,000 members. 



CHAPTER II 
MATTER AND ITS CLASSIFICATION 

Definition of matter. Since the chemist is interested primarily 
in matter and the changes it undergoes, it is important for us 
to get a clear idea of just what we mean by matter. For our 
purposes we may define matter as anything that has weight or 
occupies space. This not only includes solid and liquid materials 
that we can see and touch but also the various gases, such 
as those that make up the atmosphere and which, although in- 
visible, are very real and fit the definition of matter just given. 

Classification of matter. From the standpoint of its physical 
state it is customary to classify matter under three headings : 
namely, solids, liquids, and gases. The chemist, however, is 
interested in matter from the standpoint of its composition, 
and from this standpoint it may be classified under two 
headings : elements and compounds. 

Elements. There are a number of different substances, some 
of them well known to all of us, — such as iron, copper, and 
gold, — that have resisted all efforts to decompose them into 
simpler substances. On this account they are called elementary 
substances or, more briefly, elements. We may therefore de- 
fine an element as a substance which cannot be decomposed into 
simpler substances. About ninety such elements are known. 

It is not always easy to prove that a given substance is really 
an element. Some way as yet untried may be successful in decom- 
posing it into other simpler forms of matter. Water, lime, and 
many other familiar substances were at one time thought to be 
elements, but are now known to contain two or more elements. 
Most of the elements are solids, a few are gases, and only two — 
bromine and mercury — are liquids under ordinary conditions. 

6 



MATTER AND ITS CLASSIFICATION 7 

Compounds. On the other hand, there are many thousands 
of substances that are made up of two or more elements. 
Thus, if a current of electricity is passed through water 
(Fig. 9), the water is decomposed into two elements (both of 
which are invisible gases), known respectively as oxygen and 
hydrogen. Moreover, if we cause these two elements to com- 
bine, water is formed. Again, the ordinary iron ore known as 
hematite can be shown to consist of the two elements iron and 
oxygen. The sugar with which we sweeten our food is com- 
posed of the three elements carbon, hydrogen, and oxygen. 
Compounds, however, are not only made up of two or more ele- 
ments but each compound has a perfectly definite composition. 
For example, we shall see later that water is made up of 
hydrogen and oxygen in the ratio of one part by weight of 
hydrogen to 7.94 parts by weight of oxygen. It makes no 
difference what the source of the water is, provided only that 
it is pure ; the composition is always the same. This constancy 
of composition is characteristic of all compounds. We may 
therefore define a compound as a substance made up of two or 
more elements combined in definite proportions by weight. We 
shall learn of other characteristics of compounds as we proceed. 

Illustration. We might, in a very general way, compare the ele- 
ments to the letters of the alphabet. Just as the printed matter 
on this page is made up of single letters and of combinations of 
letters to form words, so matter is made up of elements and. of 
combinations of elements in the form of compounds. Just as there 
are a great many more words than letters, so there are a great 
many more compounds thau elements. 

Chemical changes — chemical action. We may best illustrate 
the meaning of these terms by a simple illustration. If we 
place in a test tube one or two grams of the red solid substance 
known as mercuric oxide and heat it (Fig. 5), we find that 
the color gradually changes and that there is left in the tube 
little drops of a silverlike liquid which we recognize as mercury, 



8 



CHEMISTKY AND ITS USES 




or quicksilver, as it is often called, which is used in thermometers 
and barometers. If during the heating we thrust into the mouth 
of the tube a glowing splint, the wood will burst into a flame. 
Experiments show that this bursting into flame is due to the 
presence of the invisible gas, oxygen, which is evolved on heating 

the mercuric oxide. The red solid, 
mercuric oxide, then has been de- 
composed by the heat into two ele- 
ments, the one a silverlike liquid 
and the other an invisible gas. A 
change like this is known as a 
chemical change, and in describing 
it we say that chemical action has 
taken place. Such changes are tak- 
ing place all about us, and they 
are the ones in which the chemist 
is particularly interested. Thus, 
coal and wood burn, being changed 
in the process into ashes and invisible gases. The food we eat 
is changed into the tissues of the body. In all such changes 
the substances resulting from the chemical action differ in com- 
position from the substances originally present and usually differ 
from them in appearance as well. We shall see later on that 
there are other important changes which always accompany 
chemical action. At present, for our purposes, we may define 
a chemical change as one that is attended by a change in the 
composition of matter. 

Changes that are not attended by a change in the compo- 
sition of matter, such as the breaking of a piece of glass or the 
powdering of a lump of coal, are known as physical changes. 

It follows from the statements made above that the appear- 
ance of a compound gives no clue as to what elements are 
present in it. Thus the red solid, mercuric oxide, is formed by 
the union of the silverlike liquid, mercury, with the invisible 



Fig. 5. The decomposition of 

mercuric oxide into mercury 

and oxygen by heat 



MATTEE AND ITS CLASSIFICATION 9 

gas, oxygen. Water, a colorless liquid, is formed by the union 
of two invisible gases, oxygen and hydrogen. No one would 
ever suspect simply from appearance that sugar contains the 
black solid element carbon, and yet if we heat sugar the hy- 
drogen and oxygen with which the carbon is combined in the 
sugar are expelled and the black carbon remains. 

Chemical affinity. The force that causes elements to unite and 
holds them in combination in compounds is called chemical affinity. 
We know very little about the nature of this force, just as we 
know very little about the force of gravitation. It is evident, 
however, that there is such a force, and it is convenient to 
have a name by which we can refer to it. 

Number of elements. The number of substances now con- 
sidered to be elements is not large — about ninety in all. Many 
of these are rare, and in some cases not more than a few grams 
have been obtained pure. Clarke makes the following estimate 
of the composition of the solid portion of the earth's crust : 

COMPOSITION OF THE EARTH'S CRUST 

Oxygen 47.33% Magnesium 2.24% 

Silicon 27.74% Sodium 2.46% 

Aluminium .... 7.85% Potassium 2.46% 

Iron 4.50% Hydrogen 0.22% 

Calcium 3.47% Other elements . . . 1.73% 

A complete list of the elements is given on the back cover 
page. It is not necessary to study more than one third of 
the total number of elements to gain a very good knowledge 
of chemistry. 

Elements in the human body. Comparatively few of the 
elements appear to be essential to life. The following table, 
compiled by Sherman, gives the average composition of the 
human body. So far as we can judge, these are the only ones 
upon which living organisms are dependent, though traces of 
others may be necessary. 



10 



CHEMISTRY AND ITS USES 



AVERAGE COMPOSITION OF THE HUMAN BODY 



Oxygen 


65.00% 


Phosphorus 


1.00% 


Magnesium . 


0.05% 


Carbon . 


18.00%' 


Potassium . 


0.35% 


Iron . . . 


0.004% 


Hydrogen 


10.00% 


Sulfur . . 


0.25% 


Iodine . . 


traces 


Nitrogen 


3.00% 


Sodium . . 


0.15% 


Fluorine . . 


traces 


Calcium 


2.00% 


Chlorine 


0.15% 


Silicon . . 


traces 



Occurrence of the elements. Most of the elements occur in 
nature not as uncombined substances but in the form of 
chemical compounds. When an element does occur uncom- 
bined, as is the case with gold and sulfur, we say that it 
occurs in the free state, or native ; when it is combined with 
other substances in the form of compounds, we say that it 
occurs in the combined state, or in combination. The elements 
present in our bodies are all in the form of compounds, of 
which water is the most abundant. 

Names of elements. The names given to the elements have 
been selected in a great many different ways. Some names, 
such as iron and gold, are very old, and their original meaning 
is obscure. Many names indicate some striking property of 
the element. The name bromine, for example, means " stench," 
referring to the extremely unpleasant odor of the substance. 
Other elements are named from countries or localities, as 
germanium and scandium. Still others are named from some 
mythological character, as thorium and tantalum. 

Symbols. In indicating the elements chemists have adopted 
a system of abbreviations. These are known as symbols, each 
element having a distinctive symbol. Sometimes the initial 
letter of the name is adopted to indicate the element. Thus, 
I stands for iodine, C for carbon. Usually it is necessary to 
add some other characteristic letter to the symbol, since sev- 
eral names may begin with the same letter. Thus, C stands 
for carbon, CI for chlorine, Cd for cadmium. Sometimes the 
symbol is an abbreviation of the name in some other language. 



MATTER AND ITS CLASSIFICATION 11 

In this way Fe (Latin, ferrwri) indicates iron, Cu (Latin, 
cuprum) indicates copper, and Ag (Latin, argentuni) indicates 
silver. The symbols will be found in the list of elements 
given ' on the back cover page. They will become familiar 
through constant use. 

The number of compounds. The number of compounds which 
have been described and which can be made when desired is 
very large, and each year many more are added to the list. 
About 200,000 are known that contain the element carbon as 
one constituent, and the total number listed in the large hand- 
books of chemistry is much larger. Fortunately it is not neces- 
sary to become familiar with any large number of these in order 
to gam an understanding of the principles of chemistry. 

Meaning of the terms mixture and substance. We have 
discussed the nature of elements and compounds and learned 
something of their characteristics. It is possible for us to 
mix intimately together many different elements and com- 
pounds without any chemical action taking place. The result- 
ing product is called a mixture. Thus, we may have a mixture 
of sand and salt or of sugar and flour. Such products differ 
from compounds in that their composition may be varied in- 
definitely ; moreover, in a typical mixture particles of differ- 
ent character may be distinguished, while hi a compound all 
particles, no matter how minute, are identical in composition 
and properties. When we wish to refer to some form of 
matter without regard to its composition, we often use the 
term substance. Thus, we might speak of an element, a 
compound, or a mixture as a substance. 

Method of study. We shall now proceed with a study of 
some of the more important elements. As a rule each ele- 
ment will be discussed under the following heads : (1) Prop- 
erties. By this word is meant all those physical characteristics 
of a substance by which we recognize it. This includes its 
state (solid, liquid, or gas), color, odor, and taste. It also 



12 CHEMISTRY AND ITS USES 

includes certain measured quantities such as weight, hard- 
ness, solubility, boiling point, freezing point. (2) Occurrence 
in nature. Under this heading will be discussed such topics 
as the forms in which the element occurs in nature, whether 
free or in the combined state, whether it is an abundant ele- 
ment or of rare occurrence. (3) Historical study. It is always 
of interest to know something concerning the discovery of the 
element and other items of historical interest in connection 
with it. (4) Preparation. As a rule the elements do not occur 
pure in nature but are mixed or combined with various other 
substances. It becomes necessary, therefore, to find out 
methods whereby the elements can be separated from other 
substances and thus obtained in a pure state. (5) Chemical 
conduct. Under this head will be described the various chemi- 
cal changes in which the element plays a part and the methods 
for producing these changes. In studying oxygen, for ex- 
ample, we shall want to know what other elements will com- 
bine with oxygen, the conditions necessary to bring about the 
combination, and the nature of the resulting products. (6) Uses. 
From a practical standpoint it is important for us to learn of 
the various uses to which the different elements are adapted. 
Method of recording temperatures. Unless otherwise in- 
dicated, all temperatures given in this book refer to the 
centigrade system. Any student not familiar with the cen- 
tigrade system of recording temperatures will find a brief 
explanation of it in the Appendix. 

EXERCISES 

1. What means have been mentioned for decomposing a compound? 

2. Define the following terms and give an example to illustrate each : 
(«) matter, (&) element, (c) compound, (d) mixture. 

3. Define the terms (a) chemical change, (b) chemical affinity. 

4. Does the fact that a substance undergoes no change on heating 
prove it to be an element ? 



MATTER AND ITS CLASSIFICATION 13 

5. Read over the list of elements, (a) With what ones are you 
familiar? (b) Is brass an element? 

6. Taking into account possible future discoveries, in what way may 
the list of elements (see Appendix) be changed ? 

7. How do you account for the fact that the Ancients were familiar 
with copper and gold ? 

8. Give three examples of chemical changes with which you are 
familiar. 

9. Can you tell from the appearance of a compound whether or not 
it contains iron ? 

10. If a substance such as a candle disappears on burning, does this 
prove that the matter in it is destroyed ? 

11. Calculate approximately the number of pounds of the more 
important elements present in your body. 

12. (a) What is the approximate weight of iron in your body? 
(b) Does the fact that the amount of iron is so small signify that it is 
unnecessary ? 

13. In round numbers, what part of the world is oxygen? 

14. Aluminium is much more abundant than iron. How do you 
account for the fact that iron is much the cheaper of the two metals? 

15. Compare the relative quantities of iron and silicon in the earth's 
crust. How do you account for the fact that everyone is familiar with 
iron and but very few with silicon ? 

16. Consult the dictionary for the derivation and significance of the 
following names of elements : (a) phosphorus, (b) hydrogen, (c) germa- 
nium, (d) columbium, (e) chlorine, (/) argon, (</) copper, (Ji) selenium, 
(i) thorium, (/) tantalum. 

17. (a) What is the symbol for gold? (&) From what word is this 
symbol derived? (Consult dictionary.) 

18. Give some typical examples of mixtures. 

19. How can you tell whether any given substance is a mixture or a 
compound ? 

20. Common alcohol boils at 78.3° C. At what temperature does it 
boil on the Fahrenheit scale ? 

21. Mercury freezes at — 38.87° C. Calculate the corresponding 
Fahrenheit temperature. 



CHAPTER III 
OXYGEN 

Introduction. Having become acquainted with a few of the 
characteristics of the class of substances called elements, we 
shall now turn to a more detailed study of two members of 
this class ; namely, oxygen and hydrogen. It is natural that we 
should begin with oxygen, since it is the most abundant of all 
elements, occurs in nature in great quantities in the elementary 
state, and plays such an important part in the familiar processes 
of burning and breathing. 

Properties of oxygen. Oxygen is one of the gases present 
in the atmosphere and is of fundamental importance, since it 
is the one which supports life. It is a colorless, tasteless, and 
odorless gas and is slightly heavier than air. One liter of oxy- 
gen weighs 1.429 g., while 1 liter of air weighs 1.2928 g. It 
is but slightly soluble in water, 100 volumes of water dissolv- 
ing only about 4 volumes of the gas at ordinary temperatures 
and pressures. 

It is well known that gases expand on heating, also that the 
quantity of gas that occupies a certain volume (say, that of a given 
automobile tire) may be greatly increased by pressure. It is evi- 
dent, therefore, that the weight of (say) 1 liter of a gas will vary 
with the temperature of the gas and the pressure to which it is 
subjected when weighed. We shall find later that in making such a 
statement as " 1 liter of oxygen weighs 1.429 g." it is understood 
that this is the weight of 1 liter of oxygen measured at 0° and a 
pressure of 1 atmosphere ; that is, 1033 g. per square centimeter. 

Occurrence. Oxygen is by far the most abundant of the ele- 
ments. In the free state it forms a considerable part of the 

14 



OXYGEN 15 

atmosphere. One hundred volumes of dry air contains about 
21 volumes of oxygen mixed with 78 volumes of the gas nitro- 
gen and 1 volume of other gases. Combined with other ele- 
ments, oxygen forms eight ninths by weight of water, nearly 
one half of the rocks constituting the earth's crust, and over one 
half of animal and vegetable organisms ; for example, 65 per 
cent by weight of the human body is oxygen. 

Historical. The Englishman Joseph Priestley (Fig. 6) is 
commonly regarded as the discoverer of oxygen, although 
other investigators, especially the Swedish chemist Scheele, 
had obtained it before Priestley, but had failed to attract 
attention to their discovery. The name oxygen signifies "acid 
producer." It was suggested by the French chemist Lavoisier 
(frontispiece) because he thought that the class of substances 
known as acids owe their characteristic properties to the pres- 
ence in them of this element. We now know that Lavoisier 
was mistaken in this view, for there are many acids that 
contain no oxygen. 

Priestley was born near Leeds, England, in 1733. He was edu- 
cated for the ministry, but became interested in science and spent 
much of his spare time in performing experiments. In 1774, while 
studying gases, or " airs," as he called them, he heated the com- 
pound now known as mercuric oxide by means of a large burning- 
glass and found that a colorless gas was evolved. This gas, which 
Lavoisier later named oxygen, attracted his attention because a 
candle burned in it with a brilliant flame. Later Priestley had to 
leave England because of his liberal views. He came to America 
in 1794 and settled in Northumberland, Pennsylvania, where he 
resided until his death in 1804. The house in which he lived still 
stands and is preserved as a memorial to him. 

Preparation. While oxygen is the most abundant of all the 
elements it does not occur in nature in a pure state, but always 
either mixed with other gases, as in the air, or combined with 
other elements in the form of compounds. Since it is so 



16 CHEMISTRY AND ITS USES 

abundant in the air one would naturally try to obtain it 
from this source. The separation of the oxygen from the 
other gases present in the air, however, is a difficult matter, 
requiring expensive machinery, and is only practicable when 
large volumes of the gas are desired. In the laboratory, 
where only small amounts are wanted for a study of the proper- 
ties of the element, it is most conveniently prepared by decom- 
posing certain of its compounds, either by heat or electricity, 
and collecting the oxygen evolved. The compounds usually 
employed are (1) mercuric oxide, (2) potassium chlorate, 
and (3) water. 

1. Preparation from mercuric oxide. Mercuric oxide is a solid, 
either red or yellow (depending on its method of preparation), 
and consists of 7.4 per cent oxygen and 92.6 per cent mer- 
cury. If a small quantity of this oxide is placed in a narrow 
test tube and heated (Fig. 5), it is readily decomposed into 
its constituent elements. The change may be represented in 
the following way, in which the names of the elements com- 
posing the compound are inclosed in brackets just beneath 
the name of the compound : 

mercuric oxide >• mercury + oxygen 

rmercury - ] 
Loxygen J 

The mercury is seen to deposit on the sides of the tube, 
while the presence of oxygen is shown by the fact that a 
glowing spark on the end of a splinter of wood inserted 
into the tube bursts into a bright flame. This method is too 
expensive for ordinary use, but it is of interest because of 
its simplicity and because it is the one which Priestley used 
in preparing oxygen. 

2. Preparation from potassium chlorate (usual laboratory 
method). Potassium chlorate is a white solid which has been 
found to consist of 31.9 per cent potassium, 28.9 per cent 
chlorine, and 39.2 per cent oxygen. When this compound is 




Fig. 6. Joseph Priestley (1733-1804) 

School-teacher, preacher, philosopher, scientist ; friend of Benjamin 
Franklin ; discoverer of oxygen ; defender of the phlogiston theory . 
He was horn in England, where he lived until he was sixty-one years 
of age ; he then came to the United States and spent the remainder 
of his life at Northumberland, Pennsylvania 



OXYGEN 17 

heated above its melting point the oxygen is given off, leaving 
a white solid compound of potassium and chlorine called potas- 
sium chloride. The change may be represented as follows : 

potassium chlorate >- potassium chloride + oxygen 

[potassium! rpotassiuml 

chlorine |_ chlorine J 

oxygen J 

The evolution of the gas becomes marked at about 400°. It 
is a remarkable fact that the rate at which the oxygen is 
evolved at any given temperature is greatly increased by the 
presence of small quantities of certain substances, notably 
manganese dioxide. By mixing such a substance with the 
chlorate it is possible, therefore, to obtain the oxygen rapidly 
at a lower temperature than would otherwise have to be em- 
ployed. As to the way in which the manganese dioxide pro- 
motes the decomposition, it may be said at once that we do 
not know. Apparently it undergoes no change during the 
reaction. Certainly it contributes no oxygen, for the weight 
of oxygen obtained is always 39.2 per cent of the weight of 
the chlorate used, irrespective of the presence of manganese 
dioxide. This is but one example of many in ivhich the rate of 
change is influenced by an apparently inactive substance. Such 
substances are called catalytic agents, or catalyzers, and we shall 
meet with them frequently in subsequent pages, since a great 
many chemical processes depend upon suitable catalyzers for 
their success. 

Directions for preparing oxygen. A convenient way of preparing 
oxygen from potassium chlorate is illustrated in the diagram' 
(Fig. 8) on page 18. A mixture consisting of four parts of potas- 
sium chlorate and one part of manganese dioxide is placed in the 
flask A and gently heated. The oxygen is evolved and escapes 
through the tube B. It is collected by bringing over the end of 
the delivery tube the mouth of a bottle or cylinder C completely 
filled with water and inverted in a vessel of water as shown in 
the figure. The gas rises in the bottle and displaces the water. 



18 



CHEMISTRY AND ITS USES 




Fig. 8. Preparation of oxygen from potassium 
chlorate, and method of collecting the gas 



Collection of gases. The method just described for collecting 
oxygen illustrates the general way in which gases are trans- 
ferred from one vessel to another when 
they are insoluble in water or nearly 
so. The vessel D (Fig. 8), containing the 
water in which the bottles are inverted, 
is called a pneumatic trough. 

3. Preparation from 
water. As we shall 
later see, water is a 
compound and con- 
tains 88.81 per cent 
oxygen and 11.19 per 
cent hydrogen. It is 
not practicable to de- 
compose it into its elements by heat, but the decomposition 
is easily effected by the use of electrical 
energy in the following way: 

Two tubes, A and B (Fig. 9), are filled 
with water and inverted in a vessel of 
water to which a little sulfuric acid has 
been added. A piece of platinum foil, C 
and D, attached to a wire is then brought 
under the end of each tube. When these 
wires are connected with a suitable source 
of current, supplying from 6 to 10 volts, 
bubbles of gas collect in each tube. These 
gases are oxygen and hydrogen. The 
volume of the hydrogen (tube A) liberated 
is approximately twice that of the oxygen 
(tube B). The reasons for adding sulfuric 
acid will be explained later on. 

The change tvhich takes place when 
a current of electricity is passed through 
a liquid is called electrolysis. In the above experiment the 
oxygen is said to be obtained by the electrolysis of water. 




Fig. 9. The decompo- 
sition of water into oxy- 
gen and hydrogen by the 
electric current 



OXYGEX 19 

Commercial preparation of oxygen. Oxygen is now a 
common article of commerce and can be purchased on the 
market, stored in strong steel cylinders (Fig. 10). When 
wanted in large quantities it is prepared almost entirely 
from air. The method used will be described later. 




Fig. 10. Oxygen stored in steel cylinders, ready for the market 

Laboratory methods and commercial methods. As we go along we 
shall see that the methods used in making various substances in 
the laboratory are often different from those employed commer- 
cially. In the laboratory, where relatively small quantities are 
desired, the easiest or most instructive way is preferred. In com- 
merce economy is the deciding factor. Moreover, it often happens 
that a method which will not work well on a small scale works 
admirably with commercial quantities, or that the value of a second 
product (by-product^ obtained at the same time makes a method 
a success. 

Chemical conduct. At ordinary temperatures oxygen is not 
very active. Most substances are either not affected by it or 
the action is so slow as to escape notice. At higher tempera- 
tures, however, it is very active. This may be shown by 
heating various substances in the air until just ignited and 
then bringing them into vessels containing pure oxygen, when 
they burn with increased brilliancy. Thus a glowing splint 
introduced into a bottle of pure oxygen bursts into name. 



20 



CHEMISTRY AND ITS USES 




Sulfur burns in air with a very little flame ; in oxygen the 
flame is increased in size and brilliancy (Fig. 11). Substances 
which burn in air only with great difficulty, such as iron, burn 
readily in oxygen ; while others, such as phosphorus, which 
burn readily in air, burn in oxygen with dazzling brilliancy. 
Changes which take place in burning. 
Since burning is such a common phenom- 
enon it has long attracted attention, and 
there has been much speculation not only 
as to the changes which take place in 
burning but also as to why certain sub- 
stances burn and others do not. For some 
time previous to the discovery of oxygen 
it was thought that all substances which 
burn contain in them an invisible material 
called phlogiston and that the burning of a 
substance was due to the escape of phlo- 
giston present. After Priestley discovered 
oxygen Lavoisier concluded that since 
substances burn so readily in oxygen this gas must have 
something to do with burning. This view would satisfactorily 
account for the fact that substances burn more readily in pure 
oxygen than in the air, which is only about one fifth oxygen. 
In place of merely speculating about it Lavoisier put his views 
to the test of experiments (frontispiece), which is the only 
way to find out the facts. He soon found that metals when 
burned in a vessel filled with air or oxygen increase in weight 
and that at the same time a certain amount of the oxygen 
present disappears. Moreover, he found that the increase in 
the weight of the metal in burning was equal to the weight 
of the oxygen which disappeared. He therefore came to the 
conclusion that when a metal or, indeed, any element burns it 
combines with oxygen, and that the product formed is a com- 
pound of the element and oxygen. This discovery of Lavoisier, 



Fig. 11. Burning sul 
fur in oxygen 



OXYGEN 



21 



which led to the true explanation of burning, is regarded as one 
of the greatest of all discoveries made in the field of chemistry. 

Not only do many of the elements burn in oxygen, but 
many compounds also burn in this gas. In such cases, as a 
rule, the compound is decomposed, and each of its constituent 
elements combines with the oxygen. 

Disappearance of some substances in burning. When some 
substances, such as a candle, burn they entirely disappear, 
and it might appear at first thought that this is not in 
accord with Lavoisier's explana- 
tion. A little reflection, however, 
might suggest that such sub- 
stances disappear in burning be- 
cause the compounds formed on 
burning are invisible gases which 
pass off unnoticed into the air. 
That this is the case and that 
these gases weigh more than the 
original candle is easily shown in 
the following way : 

A candle (^4) is placed on one 
pan of a balance (Fig. 12). Over it 
is suspended a wide glass tube (B) 
loosely filled with pieces of quick- 
lime or caustic potash. The whole 

apparatus is carefully counterpoised by weights (C) ; then the 
candle is lighted. As the candle burns, the pan upon which it rests 
gradually sinks, indicating that the gases formed during burning 
and absorbed in the tube over the flame are heavier than the part 
of the candle which has burned. We shall find later that the 
gases formed in the burning of the candle are the compounds 
known as carbon dioxide and water. The glass tube is filled with 
quicklime, because this substance readily absorbs both these gases. 

Oxidation ; combustion. When a substance combines with 
oxygen we say that the substance undergoes oxidation. If the 




Fig. 12. Experiment to show 
that invisible gases are formed 
when a candle burns and that 
these gases weigh more than the 
burned part of the candle 



22 CHEMISTRY AND ITS USES 

change takes place slowly no light is evolved, and unless care- 
ful measurements are made no heat is noticed. The decay of 
vegetable matter such as wood and leaves and the rusting of 
certain metals are examples of this slow oxidation. If, on the 
other hand, the process takes place rapidly, light is given off, 
and we say that the substance burns or undergoes combustion. 
These statements may be summed up as follows : Oxidation is 
the term applied to the change which takes place when a substance 
combines with oxygen. Oxidation which takes place so rapidly as 
to evolve light is called combustion. 

The term combustion as commonly used implies that oxygen is one 
of the substances entering into the change. This is always the case 
when a substance undergoes combustion in air. In a broader sense, 
however, the term is often applied to the change which takes place 
when any two substances combine with the evolution of light. 

Oxides ; products of combustion. We have seen that when 
an element burns there is formed a compound of the element 
with oxygen. Such compounds are called oxides. Thus, iron 
on burning forms oxide of iron, and sulfur forms oxide of 
sulfur. A candle is composed chiefly of carbon and hydrogen. 
When it burns the carbon unites with oxygen to form oxide 
of carbon, while the hydrogen unites with oxygen to form 
water, which is an oxide of hydrogen. All but about half 
a dozen elements form oxides, and many of them form more 
than one, so that a large number of these compounds is known. 
Some of these are solid bodies, as in the case of oxides of mer- 
cury and iron ; others are liquids, of which class water is the 
most familiar example ; quite a number are gases, as is true 
of the oxides of carbon and of sulfur. 

The oxides formed when any substance undergoes combus- 
tion are known as the products of combustion of that substance. 
Thus, oxide of iron is the product of combustion of iron, while 
carbon dioxide and water are the products of combustion of 
the candle. 



OXYGEN 23 

To sum up these facts we may say then that an oxide 
is a compound formed by the union of any one element with 
oxygen. The 'particular oxides formed when a substance under- 
goes combustion are known as the products of combustion of 
that substance. 

Rapidity (or speed) of combustion. It is a matter of common 
knowledge that the same substance may burn with different 
degrees of rapidity, and it is important for us to know how 
the rapidity, or speed, of the combustion can be increased or 
decreased. A number of factors may affect this change of 
speed, but the two most important are the following: 

1. Temperature. Everyone knows that a lump of coal if 
ignited and placed on a piece of hot iron will burn more 
rapidly than a similar one placed on a piece of cold iron. This 
is due to the fact that the hot iron rapidly increases the tem- 
perature of the coal, and this increase in temperature causes 
a corresponding increase in combustion. Within certain limits 
it may be stated in general terms that combustion takes place 
more rapidly the higher the temperature. Oxidations which may 
require years for completion at a low temperature may take 
place in a few seconds at a very high temperature. 

2. Concentration. Since ordinary combustion is due to the 
union of one or more elements with oxygen, it follows that the 
more oxygen that can be brought into contact with the sub- 
stance undergoing combustion, the more rapidly the combi- 
nation may take place. Thus a log of wood burns slowly. 
Cut the log into pieces, and the speed of combustion increases 
because more oxygen comes in contact with the wood. Split 
the pieces into fine splinters and the speed of combustion is 
still further increased. Grind the splinters into fine dust and 
float these in the air so that each little particle is surrounded 
by the oxygen in the air, and combustion, if started, will take 
place so rapidly as to be nearly instantaneous, producing a 
violent explosion. In this way we account for the explosions of 



24 



CHEMISTRY AND ITS USES 



flour mills and grain elevators, the air of which often contains 
large quantities of dust particles (Fig. 13). 

Heat of oxidation and combustion. Evidently a given sub- 
stance may either undergo a slow oxidation or it may undergo 
combustion. Thus a piece of phosphorus exposed to the air 
in a cold room slowly wastes away until it has all disappeared 




Fig. 13. That a mixture of air and dust may explode violently is shown by 

the above picture of a flour mill in Beatrice, Nebraska, taken just after a 

dust explosion had occurred 

into smoke consisting of an oxide of phosphorus, but if it is 
touched with a lighted match it takes fire and burns very rap- 
idly, giving out much heat in its combustion. The product is 
the same in both cases ; namely, oxide of phosphorus. Appar- 
ently the difference lies in the amount of heat given off, but 
very accurate experiments demonstrate that this, too* is exactly 
the same. In the one case the action is so slow that the heat 
is conducted away as fast as it is liberated, and so it escapes 
notice ; in the other it is given off so rapidly as to be very 



OXYGEN 



25 



striking. A similar relation has been found to hold true in all 
cases of combustion. The heat given off in oxidation is exactly 
the same whether the action is fast or slow, provided the same 
compound is formed. 

Spontaneous combustion. It has been found that the rate at 
which oxidation goes on is greatly increased by raising the tem- 
perature of the material undergoing oxidation. Consequently, if 
the conditions surrounding oxidation are such that the heat given 




Fig. 14. Sewage-disposal plant at Columbus, Ohio, in which the sewage is 
sprayed into the air to secure its oxidation 

off cannot escape, the temperature will steadily rise, and because 
of this the rate of oxidation will increase. The increased heat thus 
set free will still further raise the temperature, until the oxidation 
passes into active combustion, the point at which this occurs being 
called the kindling temperature. Materials taking fire in this way 
are said to undergo spontaneous combustion. It will be seen that 
the essential conditions are (1) an existing slow oxidation and 
(2) good heat insulation. Linseed oil, used in paints, undergoes 
rather rapid oxidation in air, and oily rags left by painters not 
infrequently occasion disastrous fires. Fine, dry coal in the center 
of a heap or in the closed hold of a vessel sometimes takes fire. 
Almost any finely divided combustible material, such as sawdust or 
flour, is dangerous when stored in a warm, dry place. Sometimes 



26 



CHEMISTRY AND ITS USES 



the heat of fermentation, which is a kind of oxidation, will start a 
fire in a haystack or barn if the hay is not well dried before storing. 

Importance of oxygen. Oxygen is one of the most important 
of the elements. It is essential to all forms of life except certain 
low forms of plant life. In the presence of certain minute 

microorganisms, which in some 
way assist in the process, the 
oxygen in the air acts upon 
the dead products of animal 
and vegetable life and converts 
them into harmless substances. 
In this way it acts as a puri- 
fying agent. For example, in 
sewage-disposal plants, sewage 
is forced into the atmosphere 
in fine sprays (Fig. 14) so that 
the oxygen can come in contact 
with the putrid matter in the 
sewage, thus purifying the sew- 
age and preventing it from be- 
coming a menace to health. 
Pure oxygen is used in the 
treatment of certain diseases in 
which the patient is unable to 
inhale sufficient air to supply 
the necessary quantity of oxygen. Aviators are supplied with 
the pure gas for use at high altitudes (Fig. 15). The gas is 
also used as a source of intense heat as well as in the man- 
ufacture of certain compounds. Thus large quantities were 
used during the World War in the preparation of the poi- 
son gas known as phosgene, which is a compound of oxygen, 
chlorine, and carbon. 

Solution of chemical problems. We have seen that every 
pure compound has a perfectly definite composition, no matter 




Fig. 15. An aviator fitted with an 

oxygen container and a device for 

breathing pure oxygen 



OXYGEN 27 

what may be the source of the compound. The only way to 
determine the composition of any compound is by experiment, 
but having once determined the composition of any individual 
compound in the laboratory it is possible to solve certain prob- 
lems in which the compound is involved. Thus, suppose we 
wish to know what weight of potassium chlorate is necessary 
for the preparation of 100 g. of oxygen. As already stated 
(p. 16), the composition of potassium chlorate as determined 
by experiment is as follows : potassium, 31.9 per cent; chlorine, 
28.9 per cent; oxygen, 39.2 per cent. In other words, 100 g. 
of potassium chlorate contains 39.2 g. of oxygen. To prepare 
1 g. of oxygen would require, therefore, 100 -f- 39.2, or 2.55 g., 
of potassium chlorate. To prepare 100 g. of oxygen then 
would require 2.55 x 100, or 255 g., of potassium chlorate. 

It must be kept in mind that the composition of a compound is 
always expressed in percentage by iceight. If we wish to calculate 
what weight of potassium chlorate is necessary for the prepara- 
tion of, say, 100 liters of oxygen, it is first necessary to calculate 
the weight of the 100 liters of the gas. By referring to the Appen- 
dix it will be found that 1 liter of oxygen weighs 1.429 g., hence 
100 liters of oxygen weighs 100 x 1.429, or 142.9 g. It is then 
easily possible by the method described above to calculate the 
weight of potassium chlorate necessary to prepare the 142.9 g. 
of oxygen. 

EXERCISES 

1. From your everyday experiences state how it is possible to 
change the volume of a gas. 

2. The statement is made in the text that 1 liter of oxygen weighs 
1.429 g. at a temperature of 0° and under a pressure of 1 atmosphere. 
"Why is it necessary to say anything about the temperature and pressure ? 

3. Give the approximate weight of oxygen in your body. 

4. («) In Fig. 8 why does the water stay in the inverted cylinder? 
(b) Why does the oxygen displace it? (c) When a little oxygen has 
entered the cylinder why does not all the water run out? 

5. Give the derivation of the following words : (a) pneumatic, 
(6) catalysis, (c) phlogiston. (See dictionary.) 



28 CHEMISTRY AND ITS USES 

6. Can you tell from the properties of a compound whether or not 
it contains oxygen? 

7. Neither sulfur nor oxygen has any odor. How do you account 
for the odor noticeable when the sulfur burns in oxygen? 

8. Suggest a simple way of proving that water is formed when a 
candle burns. 

9. (a) How do you account for the fact that when a lamp is lighted 
a film of moisture is deposited on the chimney ? (b) Is the film deposited 
on the inside or the outside of the chimney? (c) Why does the film 
soon disappear? 

10. When ice water is poured into a goblet, moisture often collects 
on the outside of the goblet. What is the source of this moisture ? 

11. What simple facts or experiments disprove the phlogiston theory? 

12. How would you collect a gas that is soluble in water? 

13. Since oxygen is so active, why is the free element present in the 
atmosphere in such large quantities ? 

14. Why do certain substances form ashes on burning, while others 
do not? 

15. Can you suggest a reason why some coals when burned leave 
more ashes than others? 

16. Can combustion take place without the emission of light? 

17. Is the evolution of light always produced by combustion? 

18. Suggest a simple test for pure oxygen. 

19. Distinguish between the terms (a) oxide, (&) oxidation, (c) com- 
bustion, (d) products of combustion. 

20. How can you account for the explosion of fuel gas ? 

21. Machines used for threshing wheat have been known to explode 
with great violence when in operation. Suggest a cause for the explosion. 

22. Why do substances burn more readily in pure oxygen than in air ? 

23. Why are oily rags more likely to start a fire than oil spilled on 
the floor? 

24. For what purposes is oxygen used in automobile repair shops ? 

25. Calculate the weight of 50 liters of oxygen. 

26. Calculate the volume of 50 g. of oxygen (unless otherwise noted 
it will be understood that all volumes will be measured at 0° and under 
a pressure of 1 atmosphere). 



OXYGEN 29 

27. How large a bottle measured in cubic centimeters will 1 g. of 
oxygen fill? 

28. State how many grams of oxygen are present in 100 g. of each of 
the following compounds : (a) mercuric oxide ; (6) potassium chlorate ; 
(c) water. (Calculate from percentages of these compounds given in 
the text.) 

29. (a) 1 g. of potassium chlorate will yield on heating what weight 
of oxygen? (b) What volume in cubic centimeters will this weight of 
oxygen occupy? 

30. An aviator wishes to take with him in his flight 100 liters of 
oxygen. What weight of potassium chlorate will be required for the 
preparation of this volume of oxygen? (Use the data obtained in 
problem 29 for the calculation.) 

31. A student wishes to fill five bottles with oxygen, each bottle hold- 
ing 500 cc. What weight of potassium chlorate will be required to 
prepare this volume of oxygen ? 

32. Suppose you heated 100 g. of potassium chlorate in a flask until 
all the oxygen is evolved, (a) What compound remains in the flask ? 
(&) Calculate the weight of this compound. 

33. A student prepared oxygen by heating a mixture of 10 g. of 
potassium chlorate and 2 g. of manganese dioxide, (a) What weight of 
oxygen was obtained? (b) What compounds remained in the flask? 
(c) Calculate the weight of each. 

34. What is the approximate volume of 100 g. of water (see discus- 
sion of metric system in Appendix). Suppose you decompose this water 
by electrolysis, (a) What volume of oxygen would be obtained? 
(&) Compare the volume of the water with that of the oxygen obtained 
from it. 



CHAPTER IV 
HYDROGEN 

Introduction. A great variety of materials undergo combus- 
tion, among them being coal, wood, oils, and various gases. 
One of these combustible gases, hydrogen, is of special interest 
because it is an elementary substance. 

Properties of hydrogen. Hydrogen resembles oxygen in that 
it is a colorless, tasteless, odorless gas. It is the lightest of all 
known substances, 1 liter of the gas weighing only 0.08987 g. 
Soap bubbles blown with hydrogen rapidly rise in the air, as 
do also balloons filled with the gas. The solubility of hydro- 
gen in water is very slight, being still less than that of oxygen. 
Pure hydrogen produces no injurious results when inhaled. 
Of course one could not live in an atmosphere of the gas, 
since oxygen is essential to respiration. 

Occurrence. Free hydrogen occurs in enormous quantities 
in the gases surrounding the sun and certain other stars. In the 
combined state it is widely distributed, being a constituent of 
water as well as of all living organisms and of many of the 
products derived from them, such as wood, starch, and sugar. 
About 10 per cent of the human body is hydrogen, largely in the 
form of water. Combined with carbon it forms many com- 
pounds, which, mixed together, constitute petroleum and natural 
gas. Traces of free hydrogen also occur in the atmosphere. 

Historical. Various combustible gases have been known 
from early ages, but they were long confused with each other. 
The gas hydrogen was first clearly recognized as a distinct 
substance by the English investigator Cavendish in 1766; 

30 



HYDROGEN 31 

he obtained it in pure condition and showed it to be differ- 
ent from all other known gases. It was named hydrogen by 
Lavoisier, the word meaning " producer of water." 

Preparation. Hydrogen can be prepared in a number of 
ways, the most important of which are the following: 

1. Preparation from water. Since water contains hydrogen 
and is so abundant it is but natural that we should try to 
obtain the element from this source. It is possible to liberate 
the hydrogen in water in two general ways : (a) In the first 
place we may use the method of electrolysis already described 
(Fig. 9) and in this way obtain both hydrogen and oxy- 
gen, or (F) we may act on water with certain elements which 
will combine with the oxygen of the water and liberate the 
hydrogen. For this purpose some of the metals serve our 
purpose best. In the case of a few of the metals this change 
occurs at ordinary temperatures. Thus, if a bit of the metal 
sodium is dropped on water, an action is seen to take place at 
once, sufficient heat being set free to melt the sodium, which 
runs about on the surface of the water. The change which 
takes place consists in the substitution of one half of the 
hydrogen of the water by the sodium, and may be represented 
as follows: 

sodium -f- water y hydrogen -{- sodium hydroxide 

[hydrogen - ! [~ sodium ~] 

hydrogen hydrogen 

oxygen J L°xygen J 

The sodium hydroxide formed is a white solid which re- 
mains dissolved in the excess of undecomposed water and 
may be obtained by evaporating the solution to dryness. The 
hydrogen is evolved as a gas and may be collected by suitable 
means. 

Fig. 16 represents a simple form of apparatus used in preparing 
hydrogen by the action of sodium on water. Since the sodium is 
lighter than water, it is kept under the water by pushing a pellet 



32 



CHEMISTKY AND ITS USES 



of the metal into one end of a short piece of lead or tin pipe, the 
other end of which has been hammered until closed. The pipe 
containing the sodium is then dropped 
into a trough of water. Hydrogen is at 
once evolved and is collected by bring- 
ing over it a bottle or cylinder filled with 
water, as shown in the figure. 

Other metals, such as magnesium 
and iron, decompose water rapidly but 
only at higher temperatures. When 
steam is passed over hot iron, for ex- 
ample, the iron combines with the 
oxygen of the steam, setting free all 
the hydrogen. Experiments show that 
the change may be represented as 
follows : 




Fig. 16. The preparation 

of hydrogen by the action 

of sodium on water 



hydrogen + iron oxide 

[iron ~] 
oxygenj 



iron + water — 

t hydrogen! 
oxygen J 

The iron oxide formed is a reddish-black compound identical 
with that obtained by the combustion of iron in oxygen. 

Preparation of hydrogen from iron and steam. The apparatus used 
in the preparation of hydrogen from iron and steam is shown in 
Fig. 17. A porcelain or iron tube A, about 50 cm. in length and 
2 cm. or 3 cm. in diameter, is partly filled with fine iron wire or 
tacks and connected as shown in the figure. The tube is heated 
slowly at first, until the iron is red-hot. Steam is then conducted 
through the tube by boiling the water in the flask B. The hot iron 
combines with the oxygen in the steam, setting free the hydrogen, 
which is collected over water in C. 

2. Preparation from acids (usual laboratory method). The acids 
constitute a very important class of compounds, and we shall 
be much concerned with them later in our study. For the 
present it is only necessary to know that they all contain 
hydrogen and that certain metals when brought in contact 



: 



HYDROGEN 



33 



with them dissolve, and in the change which takes place the 
hydrogen of the acid is liberated. This constitutes a very 
convenient method for preparing hydrogen in the laboratory. 
It has been found most convenient and economical in pre- 
paring hydrogen by this method to use either zinc or iron as 
the metal and either hydrochloric acid or sulfuric acid as the 
acid. Hydrochloric acid is an aqueous solution of a gaseous 




Fig. 17. The preparation of hydrogen by passing steam over hot iron 



compound known as hydrogen chloride (which consists of 2.77 
per cent hydrogen and 97.23 per cent chlorine), while sul- 
furic acid is an aqueous solution of an oily liquid known as 
hydrogen sulfate (which consists of 2.05 per cent hydrogen, 
32.70 per cent sulfur, and 65.25 per cent oxygen). 

The changes taking place in the preparation of hydrogen 
from zinc and sulfuric acid may be represented as follows: 



zinc + sulfuric acid 



[hydrogen"] 
sulfur I 
oxygen J 



zinc sulfate + hydrogen 

[zinc "1 
sulfur J 
oxygenj 



In other words, the zinc takes the place of the hydrogen in 
sulfuric acid. The resulting compound contains zinc, sulfur, 



34 



CHEMISTRY AND ITS USES 




and oxygen and is known as zinc sulfate. This remains dis- 
solved in the water in the acid. It may be obtained in the 
form of a white solid by evaporating the liquid left after the 
metal has passed into solution. 

Directions for preparing hydrogen from acids. The preparation of 
hydrogen from acids is carried out in the laboratory as follows : 
The metal is placed in a flask or wide-mouthed 
B bottle A (Fig. 18), and the acid is added slowly 

through the funnel tube B. The metal dissolves 
in the acid, while the hydrogen which is liber- 
ated escapes through the exit tube C and is 

collected over water. 

Commercial prep- 
aration. Both of the 
general methods of 
preparation described 
above have been used 
in preparing hydro- 
gen on a large scale, 
but these have been largely displaced by cheaper methods 
recently developed. The reactions involved are too compli- 
cated to describe at the present time, but will be referred to 
in a later chapter. 

Chemical conduct. At ordinary temperatures hydrogen is 
not an active element. Under suitable conditions, however, 
it combines with a number of the elements, forming many im- 
portant compounds. Thus, hydrogen and chlorine, when mixed 
together, will combine with explosive violence if heated or if 
exposed to the sunlight. The product formed in either case is 
called hydrogen chloride. Under suitable conditions hydrogen 
combines with nitrogen to form ammonia and with sulfur to 
form the foul-smelling gas hydrogen sulfide. At ordinary tem- 
peratures hydrogen and oxygen may be mixed without action. 
If the mixture is heated to about 800°, or if a name is brought 



Fig. li 



The preparation of hydrogen by the 
action of metals on acids 



HYDROGEN 



35 



in contact with it, a violent explosion takes place. Neverthe- 
less, under proper conditions hydrogen may be made to burn 
quietly in either oxygen or air. The flame produced by the 
combination is almost colorless and is very hot. The com- 
bustion of the hydrogen is due to its union with oxygen, and 
the product of the combustion is an oxide of hydrogen. That 
this compound is water may be easily shown by experiment. 




Fig. 19. Burning hydrogen in air and collecting the product of its 
combustion (water) 

Directions for burning hydrogen. The combustion of hydrogen in 
air may be carried out safely as follows : The hydrogen is generated 
in the bottle A (Fig. 19), is dried by conducting it through the tube B 
filled with some substance (usually calcium chloride) which has a 
great attraction for moisture, and escapes through the tube C, the 
end of which is drawn out to a jet. When all the air has been 
expelled from the apparatus the hydrogen may be ignited. It then 
burns quietly, since only the small amount of it which escapes from 
the jet can come in contact with the oxygen of the air at any one 
time. By holding a cold, dry bell jar or bottle over the flame in 
the manner shown in the figure, the steam formed by the combus- 
tion of the hydrogen is condensed, water collecting in drops on 
the sides of the jar. 

Hydrogen not a supporter of combustion. While hydrogen 
is readily combustible, it is not a supporter of combustion ; in 



36 



CHEMISTRY AND ITS USES 



other words, substances will not burn in it. This may be shown 

by bringing a lighted candle supported by a stiff wire into a 
bottle or cylinder of the pure gas, as shown in 
Fig. 20. The hydrogen is ignited by the flame 
of the candle and burns at the mouth of the 
cylinder, where it comes in contact with the 
oxygen in the air. When the candle is thrust 
up into the gas its flame is extinguished. If 
slowly withdrawn, the candle is relighted as it 
passes through the layer of burning hydrogen. 
Reduction. On account of its tendency to 
combine with oxygen, hydrogen has the power 
of abstracting it from many of its compounds. 
Thus, if a stream of hydrogen generated in A 
(Fig. 21) and dried by passing through the tube 
B (filled with calcium chloride) is conducted 
through the tube C, which contains some copper 
oxide heated to a moderate temperature, the 

hydrogen takes away the oxygen from the copper oxide. The 

change may be represented as follows : 

copper oxide + hydrogen 

t copper 1 
oxygenj 

Q 



Fig. 20. Flame 
of a candle ex- 
tinguished by 
hydrogen 



-*■ water + copper 

t hydrogen! 
oxygen J 



The water formed 
collects in the cold 
portions of the tube 
C, near its end, and 
drops into the glass. 
In this experiment 
the copper oxide is 
said to undergo re- 
duction. Reduction 
may therefore he de- 
fined as the process 




Fig. 21. The reduction of hot copper oxide by 
a stream of hydrogen 



HYDROGEN 



37 



of withdrawing oxygen from a compound. As we shall see, the 
term reduction is also used with a somewhat different meaning. 
Relation of reduction to oxidation. At the same time that 
the copper oxide is reduced, it is clear that the hydrogen is 
oxidized, for it combines with 
the oxygen given up by the 
copper oxide. The two processes 
are therefore very closely related, 
and it usually happens that when 
ohe substance is oxidized an- 
other substance taking part in 
the reaction is reduced. The 
one which gives up its oxygen is 
called an oxidizing agent, while 
the other, which unites with the 
oxygen of the oxidizing agent, 
is called a reducing agent. 

The oxyhydrogen blowpipe ; the 
blast lamp. The oxyhydrogen blow- 
pipe is a device for burning pure 
hydrogen in pure oxygen. It was 
devised by the American chemist 
Eobert Hare (Fig. 22) in 1801 
and was formerly used as a source 
of intense heat. A similarly con- 
structed apparatus, known as the 

blast lamp, is now used in all laboratories (Fig. 23). It consists 
of two tubes, one inside the other. Coal gas or natural gas (both 
of which are gaseous compounds of hydrogen) is forced through 
the inner tube (Fig. 23) and lighted. Air is then forced through 
the outer tube in quantities sufficient to burn the gas. 

Uses of hydrogen. Hydrogen is used for inflating balloons 
and dirigible airships such as the Zeppelins, so largely used in 
the World War (Fig. 7). Its chief peace-time use is in the re- 
fining of certain oils, such as cottonseed oil, whereby these oils 




Fig. 22. Robert Hare (1781-1858) 

An early American chemist ; the in- 
ventor of a number of ingenious 
laboratory appliances, including the 
oxyhydrogen blowpipe 




38 CHEMISTRY AND ITS USES 

are not only purified but are transformed into solid fats which 
may be used in cooking to replace lard (see hydrogenation of 
oils). Hydrogen is also coming into use in the preparation of 

ammonia, as will be explained later. 
Further statements concerning 
the solution of chemical problems. 
Many compounds can be made by 
the direct union of the elements 
as present in the compound. Thus, 
hydrogen and oxygen unite directly 
to form water. Knowing the com- 
position of the compound, one can 
tell in what proportion the elements 
Fig. 23. The ordinary labora- wiU combine to f orm the compound. 
tory blast lamp . . 

Ihus the composition ot water, as 
determined by experiment, is hydrogen 11.19 per cent and 
oxygen 88.81 per cent ; in other words, 11.19 g. of hydrogen 
will combine with 88.81 g. of oxygen to form 100 g. of water. 
If the two gases are brought together in other proportions 
than these and ignited, combination will take place, but a 
definite amount of the gas in excess will be left uncombined. 
For example, if 11.19 g. of hydrogen were mixed with, say, 
125 g. of oxygen, only 88.81 g. of the oxygen would enter 
into combination with the hydrogen, leaving 125—88.81, or 
36.19 g., of oxygen uncombined. 

EXERCISES 

1. (a) What is the approximate weight of the hydrogen present in 
your body ? (6) In what form is most of it found ? 

2. What per cent of water, by weight, is hydrogen? 

3. (a) What element other than hydrogen did Lavoisier name? 
(/>) What was Lavoisier's greatest contribution to chemistry? 

4. Should you expect hydrogen and oxygen to occur together in a 
free state in nature ? 



HYDROGEN 39 

5. In preparing hydrogen as illustrated in Fig. 17 do you see any 
reason why the tube A should be inclined as represented in the figure ? 

6. (a) In preparing hydrogen according to Fig. 18 what is the com- 
position of the gas that at first escapes through the tube C'i (b) Is such 
a gas explosive, and, if so, under what conditions ? 

7. Contrast the properties of hydrogen and oxygen. 

8. What factors determine the choice of the method used for pre- 
paring a substance commercially ; that is, on a large scale ? 

9. In Fig. 19 why is it necessary to dry the hydrogen by means of 
calcium chloride in the tube B ? 

10. From Fig. 21 suggest a way for determining by experiment the 
weight of water formed in the reaction. 

11. Why does a mixture of hydrogen and oxygen explode more 
violently than a mixture of hydrogen and air? 

12. Weight for weight, which will yield the greatest volume of 
hydrogen : hydrogen chloride, hydrogen sulfate, or water ? 

13. Contrast the terms (a) oxidation and reduction ; (£>) oxidizing 
agent and reducing agent. 

14. Why does the oxyhydrogen blowpipe give a hotter flame than 
the blast lamp ? 

15. (a) Would the relative proportions of hydrogen and oxygen 
used in an oxyhydrogen blowpipe make any difference in the in- 
tensity of heat obtained? (b) If so, what proportions would give the 
best results? 

16. What is a serious objection to the use of hydrogen for inflating 
balloons used for war purposes ? 

17. How much heavier is oxygen than hydrogen? 

18. Calculate the weight of 50 liters of hydrogen. 

19. Calculate the volume of 50 g. of hydrogen. 

20. How large a bottle, measured in cubic centimeters, will 1 g. of 
hydrogen fill ? (Compare with exercise 27 of the preceding chapter.) 

21. (a) What weight of hydrogen is present in 100 g. of hydrogen 
chloride ? (b) If liberated, what volume would this occupy ? 

22. One gram of hydrogen chloride will yield (a) what weight of 
hydrogen ? (b) what volume of hydrogen ? 

23. One gram of hydrogen sulfate will yield (a) what weight of 
hydrogen ? (b) what volume of hydrogen ? 



40 CHEMISTRY AND ITS USES 

24. A student wishes to fill five bottles with hydrogen, each hold- 
ing 1 liter. What weight of hydrogen sulfate will be required for its 
preparation ? 

25. The dirigible R-34 (see Fig. 7) has a capacity of 2,000,000 cubic 
feet. One cubic foot equals 28.32 liters. Calculate the weight in kilo- 
grams of hydrogen sulfate required to prepare sufficient hydrogen to 
fill the dirigible. 

26. One gram of hydrogen is mixed with 20 g. of oxygen and the 
mixture is exploded. Calculate the weight of water formed. 

27. Hydrogen and chlorine unite to form the gas known as hydrogen 
chloride, (a) 100 g. of hydrogen will combine with how many grams 
of chlorine (p. 33) ? (6) How many grams of hydrogen chloride will 
be formed? 

28. Suppose that 10 1. of hydrogen were passed over copper oxide 
(Fig. 21) and that all the hydrogen combined with oxygen in the copper 
oxide to form water. What weight of water would be formed ? 

29. Suppose that 10 g. of water were decomposed by electrolysis 
(Fig. 9). (a) What weight of hydrogen and oxygen respectively would 
be formed (p. 18) ? (b) What volumes would the hydrogen and oxygen 
occupy ? 



CHAPTER V 



HOW GASES ACT; HOW THEY ARE MADE UP 



Introduction. It will be remembered that iii describing the 
properties of oxygen and hydrogen the weight of one liter of 
each gas was given. A moment's reflection will make it clear 
that these weights must refer to 
some set of arbitrary conditions, 
for it is a familiar fact that the 
volume of a given quantity of 
a gas varies both with changes 
in pressure and with changes in 
temperature. Consequently the 
weight of a liter of a gas must 
also be variable. 

Variation of volume with pres- 
sure : law of Boyle. That the 
volume occupied by a given 
weight of a gas can be altered 
by changing the pressure is 
familiar to everyone who has 
pumped air into a bicycle or 
automobile tire. As early as 
1660 Robert Boyle, an English 
investigator (Fig. 24), reached the following conclusion, known 
as Boyle's law : The volume occupied by a given iveight of a gas 
is inversely proportional to the pressure, provided the temperature 
remains constant. Thus, if a given weight of a gas occupies 
a volume of 1000 cc. when subjected to a certain pressure, it 
will occupy a volume of 500 cc. if the pressure is doubled, or 

41 




Fig. 24. Robert Boyle (1627-1691) 

One of the most accurate of the 

early experimenters in chemistry 

and physics 



42 CHEMISTRY AND ITS USES 

of 2000 cc. if the pressure is diminished to one half. This 
means that for a given weight of a gas the product of the 
pressure into the volume will remain constant, no matter how 
either one may be altered. Designating the pressure and vol- 
ume under one set of conditions by P and V, and under a 
different set by P 1 and V^ Boyle's law may be stated thus: 

PV=P 1 V 1 

It is a remarkable fact that all gases act in this same way. 
Standard pressure. For practical purposes, therefore, we 
must choose some standard pressure under which we will 
agree to measure all gas volumes. This is most conveniently 
chosen as the average pressure of the atmosphere at the 
sea level. This is equal to 1033 g. per square centimeter. 
In place of expressing the pressure in this way it is much 
more convenient to express it in terms of the height of the 
column of mercury which the pressure of the atmosphere 
will sustain. Expressed in this way the standard pressure 
is equal to that exerted by a column of mercury 760 mm. 
in height, this being the average height of the barometer at 
the sea level. 

Illustration of the law of Boyle. The following example will not 
only make the meaning of the law clear but will also show how the 
law enables us to calculate the changes in the volume of a gas due 
to changes in pressure : 

A gas measured under a pressure of 720 mm. had a volume 
of 620 cc. What volume will this gas occupy under standard 
pressure, 760 mm., the temperature remaining constant ? 

According to Boyle's law, PV = P 1 V V Substituting the 
values given in the problem, we have 760 x V— 720 x 620, or 
V = 587.4 cc. 

Variation of volume with temperature. If the pressure is 
held constant, all gases expand when the temperature is raised 
and contract, when it is lowered, and it is a remarkable fact 
that the volumes of all gases change to the same extent for 



HOW GASES ACT 



43 



Centigrade 



100 1 



-119 ' 



•252.7 
•&73 



Boiling Pointy 
of Water 



Freezing Point 
of Water 



Boiling Po%nl_ 
of Ozone 



Boiling Point 
of Hydrogen" 
— Absolute Zero- 



Absolute 



373' 



273' 



a given variation in the temperature. Let us suppose that the 
volume of a gas has been measured at 0° on the centigrade 
scale. Experiment has shown that for each degree the tem- 
perature is raised above 0°, the volume of the gas increases 
by 2T3" °^ ^he YOmme it occupied at 0°. Similarly, for each 
degree the temperature is lowered 
below 0°, the volume diminishes by 
2 j g- of the volume it occupied at 0°. 
To take a simple illustration : If the 
volume of a gas at 0° is 273 cc, then 
at 1° it will be 274 cc, while at -1° 
it will be 272 cc. At 5° it will be 
278 cc, while at - 5° it will be 268 cc. 
It is evident, then, that at this rate 
of change the volume at — 273° will 
be zero. Of course this cannot really 
happen, and experiment shows that 
before this temperature is reached, 
all gases have changed into liquids 
or solids. Helium, the most difficult 
gas to liquefy, passes into a liquid 
at - 268.7°. 

The absolute scale. For many pur- 
poses it has been found convenient 
to use a new scale of temperature 
known as the absolute scale, on which 
the divisions are of the same size as 

those on the centigrade scale, but the zero point is the same 
as — 273° on the centigrade scale ; that is, the temperature at 
which the volume of any gas would apparently become zero 
(see preceding paragraph). The 0° on the centigrade scale 
would then be 273° on the absolute scale. On such a 
scale all temperatures are above the zero point. To con- 
vert readings on the centigrade scale to the corresponding 



203° 



Fig. 25. Comparison of the 

centigrade with the absolute 

scale of temperature 



44 



CHEMISTRY AND ITS USES 



ones on the absolute it is only necessary to add 273. Thus, 
20° C. = 20 + 273°, or 293° A., while - 20°= - 20 + 273, or 
253° A. Fig. 25 gives a comparison of the centigrade and 
absolute scales at a number of temperatures. 

The law of Gay-Lussac (or of Charles). A general statement 
can now be made in regard to the effect of temperature on the 

volume of a gas: The volumes 
occupied by a given weight of 
a gas at different temperatures 
are proportional to the absolute 
temperatures, provided the pres- 

If V 




sure remains constant 
and 



V 1 are the volumes at the 



temperatures T and T v then 



F:K 1 = 



or 



T: T 



i' 



V = 



V X T 



Fig. 26. Joseph Louis Gay-Lussac 
(1778-1850) 

A distinguished French chemist who 
contributed much to our knowledge 
of gases and their combining ratios 



The above generalization is 
called the law of Gay-Lussac 
(Fig. 26) or of Charles, since it 
was formulated independently 
by these Frenchmen in 1801. 



Illustration of the law of Gay- 
Lussac. The following example 
will make the meaning of the law clear. The volume of a certain 
gas measured at a temperature of 70° is 650 cc. What will be its 
volume at 10°, the pressure remaining unchanged ? 
First reduce the centigrade readings to absolute : 

70° C. = 70 + 273 = 343° A. ; 10° C. = 10 + 273 = 283° A. 

Then substitute the appropriate values in the above equation: 

650 x 283 



V = 



343 



-, or V= 536.3 cc. 



HOW GASES ACT 45 

Variations in volume due to changes both in pressure and 
temperature. In case both pressure and temperature change, 
then the correction may be made for each in succession, as 
illustrated in the following example: 

A certain weight of gas measured 500 cc. at a temperature of 
100° when subjected to a pressure of 760 mm. Calculate the 
volume which this gas will occupy at a temperature of 50° and a 
pressure of 740 mm. 

First make the correction for pressure : 

pv = p 1 v 1 

740 x V = 760 x 500, or V = 513.5 cc. 
Next make the correction for temperature : 

y= V£ F= 6ia|x323 _ or F = 4446cc . 

Standard conditions. Since the volume of a gas varies with 
both temperature and pressure, it is essential that we select 
both a standard temperature and a standard pressure to which 
we shall agree to refer all gas volumes. We have already noted 
that the standard pressure adopted is that exerted by a column 
of mercury 760 mm. in height. As a standard temperature, the 
temperature of melting ice is chosen. This is 0° centigrade or 
273° absolute. Whenever the volume of a gas is given it is 
always assumed, unless otherwise specified, that the volume 
given is that occupied by the gas under standard conditions. 

Vapor pressure of water. As a rule gases are measured 
in the laboratory by collecting them over water in a grad- 
uated tube, as represented in Fig. 27. Before the reading is 
taken the tube A is first raised or lowered until the level of 
the water is the same within and without the tube ; the in- 
closed gas is then under atmospheric pressure. But to some 
extent water has evaporated into the tube, and a part of the 
volume inclosed is due to water vapor and not to the gas. If 
the water vapor could be removed, the gas would occupy a 



46 



CHEMISTRY AND ITS USES 



smaller volume. The downward pressure upon the water in 
the tube A is partly due to the pressure exerted by the gas 
and partly to the pressure exerted by the water vapor. If 
we could find out how much pressure the 
water vapor exerts upon the surface of 
the water within the tube and subtract 
this from the atmospheric pressure, we 
should have the pressure which the gas 
itself exerts at the volume which it occu- 
pies. The pressure due to the water vapor 
is called the vapor pressure of water. It in- 
creases steadily as the temperature rises, 
and the table in the Appendix gives its 
value over the usual range of tempera- 
tures that occurs in the laboratory. 




Fig. 27. Measuring 
a volume of gas col- 
lected over water 



Solution of problems involving calculation of 
volumes of gases collected over water. By 

referring to Fig. 27 it is evident that the 
atmosphere pressing down upon the surface 
of the water in the cylinder tends to force 
the water up in the tube A and so compresses the gas in the 
tube. The aqueous vapor mixed with the gas within the tube 
tends to push the water down and thus acts in opposition to the 
atmospheric pressure. It is evident, therefore, that to obtain the 
value of the effective pressure to which the gas in the tube is 
subjected, we must subtract the value of the vapor pressure of 
water from the value of the pressure indicated by the barometer. 
The following example will make this clear : 

A gas measured over water has a volume of 300 cc. when the 
barometer reads 740 mm. and the thermometer 20°. Calculate 
the volume which the gas will occupy if the pressure is increased to 
760 mm., the temperature remaining constant. The effective pres- 
sure on the gas equals 740 mm. less the vapor pressure of water at 
20°, which is 17.51 mm. (Appendix), or 740 - 17.51 = 722.49 mm. 
Substituting these values in the equation expressing Boyle's law, 
we have V x 760 --= 722.49 x 300 , or V = 285 cc. 




Fig. 28. Amedeo Avogadro (1776-1856) 

A celebrated Italian scientist who first advanced the fundamental 
principle of modern chemistry known as Avogadro's Principle. He 
was professor in the University of Turin, Italy. The engraving is 
from a photograph of a statue erected at Turin to his memory 




Fig. 29. Justus Liebig 
(1803-1873) 

A great German chemist and 
teacher. A pioneer chemist, es- 
pecially in the application of 
chemistry to agriculture 




Fig. 30. Michael Faraday 
(1791-1867) 

A distinguished English chemist 
and physicist. Noted especially 
for his study of electrolysis. He 
discovered a number of com- 
pounds, including benzene, from 
which most aniline dyes are 
prepared. He was the first to 
study the methods of liquefying 
gases and was very skillful in 
performing experiments 



HOW GASES ACT 47 

Nature of gases ; Avogadro' s principle. It is a surprising 
fact that all gases are equally affected by changes in temper- 
ature and pressure. This property of gases suggests that there 
must be some marked similarity in the make-up of all gases. 
The most likely guess about this that anyone could think of 
was advanced many years ago and was known as the kinetic 
theory. The main points of this theory are as follows : (1) all 
gases are made up of minute particles called molecules ; (2) all 
the molecules of any definite gas are identical in every way, 
but the molecules of different gases differ at least in weight 
and composition ; (3) the molecules are in very rapid motion, 
constantly hitting each other as well as the sides of any contain- 
ing vessel; (4) when a gas is heated, the motion of the molecules 
is increased. In accordance with this idea of a gas, the mole- 
cules do not completely fill all the space in which they are 
confined, just as the people in a room do not fill all the space 
in the room. If the gas is compressed, the molecules are not 
changed in any way, but are simply crowded more closely 
together. One may get a rough idea of the make-up of a gas 
in which the molecules are enormously magnified by noting 
the conditions that exist in a swarm of bees. The swarm is 
made up of individual bees all alike and all in rapid motion. 

The above view of matter, although originally advanced as 
a plausible guess, is now known to be true. Indeed, we can 
count the molecules and even approximate their weight. More- 
over, we know that equal volumes of all gases under the same 
conditions of temperature and pressure inclose the same number 
of molecules. This view, originally advanced by the Italian 
scientist Avogadro (Fig. 28), is now known to be a statement of 
fact and is termed Avogadro' s principle. It is a very important 
fact, and we shall refer to it in later chapters. We shall find 
later that solids and liquids also are made up of molecules,- 
but in these the molecules are closer together and not in such 
rapid motion. Avogadro' 's principle, however, applies only to gases. 



48 CHEMISTRY AND ITS USES 



EXERCISES 

Note. All temperatures are expressed in the centigrade system unless other- 
wise stated. In solving problems the student will take into account the vapor 
pressure of water only when the statement is made that the gas is collected 
over water. 

1. Give two illustrations of Boyle's law from everyday experience. 

2. Why is the bottom of a balloon left open and not tightly closed? 

3. Why does a balloon tend to fall toward evening and rise at 
midday ? 

4. Why does the carburetor on a motor car have to be adjusted 
with changes in the atmospheric conditions in order to secure the 
greatest efficiency? 

5. How could you prevent the expansion of a gas when heated ? 

6. Why is it not advisable to innate automobile tires to as high 
pressures in summer as in winter? 

7. How can you account for the explosion of a steam boiler ? 

8. What evidence can you give tending to show that the amount 
of water vapor taken up by a gas increases with the temperature? 

9. What would be the effect of relieving a gas from any pressure 
whatever ? 

10. Show that the kinetic conception of gases is in accord with the 
facts stated in the law of Boyle and the law of Gay-Lussac. 

11. A gas has a volume of 1 liter under a pressure of 730 mm. What 
will be its volume if the pressure is increased to standard (760 mm.), 
the temperature remaining constant ? 

12. In exercise 11 suppose that the gas had a volume of 1 liter meas- 
ured over water at 20°. Calculate the volume of the dry gas under 
standard pressure, the temperature remaining constant. Compare your 
answer with that obtained in exercise 11, 

13. A gas under standard pressure measured 500 cc. Calculate the 
pressure necessary to diminish the volume to 425 cc. 

14. A gas measured 200 cc. in a laboratory in which the temperature 
was 20°. During the night the temperature fell to 15°, but the barometric 
pressure remained constant. What volume did the gas then occupy ? 

15. Fifty liters of oxygen at a temperature of 20° was heated until 
the temperature was 50°, the pressure remaining constant. Calculate 
its volume at this increased temperature. 



HOW GASES ACT 49 

16. A gas measured 200 cc. in a laboratory in which the thermome- 
ter registered 20° and the barometer 740 mm. Calculate its volume 
under standard conditions. 

17. A student wishes to fill five bottles, each holding 1000 cc, with 
oxygen in a laboratory in which the temperature is 20° and the baromet- 
ric reading is 750 mm. The gas must be collected over water, (a) What 
volume will this gas occupy under standard conditions ? (b) What will 
this volume of oxygen weigh ? (c) What weight of potassium chlorate 
will be necessary to prepare this weight of oxygen? 

18. An automobile tire having a capacity of 1200 cu. in. is inflated 
to a pressure of 70 lb. per square inch. The temperature of the air at 
the time of inflation is 15°, but on rapid driving the temperature is in- 
creased to 40°. What is the pressure now exerted by the air in the_ 
tube, assuming that the volume of the tire remains constant ^/^Sugges- 
tion : First calculate the volume which 1200 cu. in. of air measured at 
15° and a pressure of 70 lb. per square inch would occupy if the tempera- 
ture were increased to 40°, the pressure remaining constant at 70 lb. ; 
then calculate the pressure necessary in pounds per square inch in order 
to decrease this volume back to 1200 cu. in., the temperature remain- 
ing at 40°. The calculation may be simplified in the following way : 
Recall that the volume of a gas varies with the pressure, also with the 
absolute temperature. In the case of the tire, however, the volume 
is constant; hence the pressure will vary directly with the absolute 

temperature : that is, 

^ ' P:P l =T:T 1 

Substitute in the above equation, and solve for P.) 



CHAPTER VI 
MATTER AND ENERGY 

Conservation of matter. We have seen that matter may be 
defined as anything that has weight or occupies space (p. 6) 
and that the chemist is interested in the various changes which 
matter undergoes (p. 1). A great many careful experiments 
have proved that while the form and the properties of matter 
change as a result of chemical action, the amount of matter 
(that is, its mass) remains constant. To put it in another way : 
no one has yet discovered any way of creating or destroying 
matter. This very important truth, first clearly recognized by 
Lavoisier, is known as the law of conservation of matter. This 
law may be stated thus: In all the changes through which a 
given quantity of matter may pass, its mass remains unchanged. 

Mass and density. While the term mass as applied to matter 
means the quantity of matter, the term density means the mass 
per unit volume. In the metric system the unit of mass is the 
gram and the unit of volume is the cubic centimeter. In these 
units the density of a substance is the number of grams in 1 cc. 
of the substance. Since 1 cc. of water weighs very nearly 1 g., 
the number stating the density of a substance also gives, for 
all practical purposes, its weight as compared with an equal 
volume of water. 

States of matter. It is well known that water may occur in 
three very different conditions, depending upon the tempera- 
ture ; namely, solid, liquid, and gaseous. These are called the 
three states of matter. This is not a peculiarity of water, for 
all substances exist in all three states, save only when the 

50 



MATTER AND ENERGY 



51 



temperature required for the melting of the solid or the vapori- 
zation of the liquid is so high that decomposition takes place 
before the change is effected. Thus we have seen that mercuric 
oxide decomposes before it melts. Potassium chlorate can be 
melted without difficulty, 
but it decomposes before 
it boils. 

Freezing and melting 
points. A solid normally 
passes into a liquid at a 
perfectly definite temper- 
ature, called its melting 
point. A given weight of 
any solid, in melting, ab- 
sorbs a definite quantity 
of heat, the exact amount 
absorbed depending upon 
the solid. This is known 
as the heat of fusion. On 
the other hand, the liquid 
formed tends to pass back 
into the solid state at this 
same temperature, called 
the freezing point, and in 
so doing the heat of fusion is given out again. In this case it 
is called the heat of solidification. For example, water freezes 
and ice melts at the same temperature ; namely, 0°. Moreover, 
1 g. of ice at 0°, in melting, absorbs a definite quantity of heat, 
while 1 g. of water at 0°, in freezing, gives out this same 
quantity of heat. 

Sometimes it is possible to cool the liquid below the freezing 
point, and it is then said to be undercooled. If a fragment of the 
solid is dropped in the undercooled liquid, solidification will at once 
begin (Fig. 31), and the temperature will rise to the true freezing. 




Fig. 31. A supercooled liquid quickly so- 
lidifies when a crystal of the solid is added 



52 CHEMISTRY AND ITS USES 

point and remain there as solidification continues. The freez- 
ing point is therefore best defined as the temperature at which the 
liquid and the solid can be mixed without change in temperature. 

Vaporization. There is no definite temperature at which a 
liquid passes into a vapor, or gas. Water exposed to the air 
vaporizes at all temperatures, and even ice and snow evaporate 
during weather in which no melting occurs. The higher the 
temperature, the more rapid the evaporation. 

Boiling point. The escape of vapor from an open vessel 
containing a liquid is hindered by the air which presses upon 
the surface of the liquid. The vapor escapes by making its 
way slowly through the air, but when the vapor pressure just 
exceeds the pressure exerted by the air, there is nothing to 
prevent the vapor from making its escape as fast as it forms, 
pushing the air before it. When this is the case any addi- 
tional heat applied to the liquid will not raise its temperature, 
but will merely increase its rate of evaporation. 

This temperature is called the boiling point of the liquid, 
and the boiling point may be defined as the temperatuxe at which 
the vapor pressure of the liquid just exceeds the opposing pressure 
of the atmosphere. It will be noticed that the boiling point of a 
liquid is not fixed, but depends upon the atmospheric pressure. 

Liquefaction of gases. At ordinary pressures water above 
100° is a gas. If it is either cooled or compressed, or both, it 
assumes the liquid state, and if cooled still further it assumes 
the solid state. This suggests that perhaps substances that are 
gases at ordinary temperature may be liquefied if cooled and 
compressed sufficiently. It is easy to show that this is true, at 
least in the case of some gases. For example, the gas known 
as sulfur dioxide assumes the liquid state when cooled to — 8° 
at ordinary pressure, or if compressed sufficiently it liquefies at 
ordinary temperatures. For a long time some of the gases, such 
as hydrogen and oxygen, resisted all efforts to liquefy them 
and were known as permanent gases. At present, however, all 



MATTER AND ENERGY 



53 



known gases have been liquefied and solidified. Liquid oxygen, 
for example, is a colorless liquid boiling at — 183°, while liquid 
hydrogen is the lightest liquid known and boils at— 252.7°. 

At present machines are used for liquefying gases. These 
consist of an engine for compressing the gas and a liquefier, in 
which the gas is finally brought to the liquid state. The lique- 
fier consists of a system of tubes and a valve so arranged that 




Fig. 32. A common type of apparatus used for liquefying air 

when a portion of the compressed gas is allowed to escape 
through the valve the heat absorbed in the expansion of the 
escaping gas is taken from the remainder of the compressed 
gas. This process is continued until the gas is cooled suffi- 
ciently to become a liquid under the pressure used. 

Eig. 32 shows a common form of apparatus used for liquefying 
air. The engine A compresses the air, the pressure being indicated 
by the gauges above B. A great deal of heat is generated in com- 
pressing the gas, and this is absorbed by cold water run through 
pipes in B. The compressed air then passes to the liquefier B, 
where it is liquefied. The liquid air flows from the liquefier into 
the Dewar flask C. 



54 



CHEMISTRY AND ITS USES 




Fig. 33. A Dewar 

flask for holding 

liquid air 



Dewar flasks ; thermos bottles. Liquid air may be kept for 

some hours in a special form of flask devised by the Scottish 
scientist Dewar, known as a Dewar flask. 
This consists of two concentric vessels 
(Fig. 33) of any convenient shape. These 
are joined together at the upper rim only, 
and the space between them is exhausted by 
an air pump. The vacuum serves as the best 
possible insulator to prevent heat conduction. 
The surface of the outer flask is often sil- 
vered in order to reflect the external heat and 
thus prevent its absorption. The so-called 
thermos bottles (Fig. 34) are constructed on 

the same plan and are very effective for keeping liquids either 

hot or cold for several hours. 

Amorphous and crystalline matter. Sometimes the particles' 

of which a piece of solid matter is composed have no definite 

form and even under. the microscope show no 

sharp edges or flat surfaces. Such solids are 

said to be amorphous. 

More often a careful examination of a solid 

will show that it is made up of a great many 

particles, each of which has sharp edges and 

flat surfaces. Such solids are said to be crys- 
talline, and each individual piece is called a 

crystal. While crystals have a great variety 

of forms (Fig. 35), yet for any given sub- 
stance the crystalline form is perfectly definite. 

Crystals range in sizes from microscopic to 

very large, a single quartz crystal found in 

California, weighing over a ton. 

Energy. We sometimes say that a certain boy possesses a 

great deal of energy, meaning thereby that he has a great 

capacity for work. We recognize this same capacity for work 



R 



Fig. 34. A ther- 
mos bottle 



MATTEK AND ENERGY 55 

in inanimate things. Thus, steam highly compressed in a boiler 
moves the railway train. The electricity present in power lines 
moves the street car. While it is difficult to frame a satisfactory 
simple definition of energy, for our purposes we may define it 
as follows : Energy is capacity for work, or the ability to do work. 
Transformation of energy. There are a great many forms of 
energy, such as the energy of heat, of electricity, of motion, and 
of light, for all these can perform work. Now it is entirely 
possible to convert each of these forms into the other forms ; 



\jjf 




Fig. 35. Some typical examples of crystals 

thus, it is the heat generated under the steam boiler that pro- 
duces the motion of the machinery. The energy of the falling 
water at Niagara Falls produces electric energy, and this in 
turn is sent through the wires to factories, and there utilized 
for many purposes, such as producing light, heat, and motion. 

Fig. 36 (p. 56) illustrates a few familiar transformations of 
energy. The heat of the flame A is converted into mechanical 
energy in the heat engine B. The motion of the engine is com- 
municated to the small dynamo C, where it is converted into mag- 
netic and electrical energy. The electrical energy is changed into 
heat and light in the incandescent lamp D and into chemical energy 
(p. 57) by the decomposition of water in E. 

Conservation of energy. While it is thus possible to convert 
one form of energy into another, all the experiments per- 
formed show that in these transformations no energy is created 
or destroyed, but that a definite quantity of energy of one kind 



56 



CHEMISTKY AND ITS USES 



always gives a definite and equivalent quantity of energy of 
another kind. This truth is known as the law of the conser- 
vation of energy and may be stated as follows : Energy may 
be changed from one form to another, but it can be neither created 
nor destroyed. 

Measurement of energy. Since changes in energy are so 
constantly taking place all about us, it is a matter of great 
practical importance to devise units for the measurement of 
energy and methods for making the measurement. In general, 




Fig. 36. Diagram illustrating some transformations of energy- 



each kind of energy must have its own units of measurement, 
just as with matter we have centimeters for lengths, liters for 
volumes, and grams for weights. In some of its forms energy 
is very difficult to measure directly. Indeed, neither units nor 
methods for the direct measurement of some forms of energy 
have as yet been devised. In such cases it is necessary to trans- 
form the energy into a form more convenient for measurement. 
In most cases it is changed into heat or electrical energy for 
this purpose, since these are easy to measure. 

Measurement of heat. A quantity of heat energy is meas- 
ured by observing to what extent it will change the tempera- 
ture of a given mass of some standard substance. Water has 



MATTER AND ENERGY 



57 



been chosen as the standard, and the unit of heat is called the 
calorie (designated by the abbreviation col.'). It is defined as 
the quantity of heat required to change the temperature of one 
gram of water one degree on the centigrade scale. 

The actual measurement of the quantity of chemical energy 
transformed into heat in any definite change is accomplished by 
the use of an apparatus called the calorimeter, represented in 
Fig. 37. The change is arranged to take place in solution in a 
measured volume of water contained in 
a thin-walled metal vessel A. This is 
placed within a double-walled vessel B, 
which contains water at the temperature 
of the room. The thermometer C indi- 
cates when the water has reached this 
temperature. This water is to prevent 
the influence of heat from without, and 
as an added precaution the vessel is cov- 
ered with a thick layer of nonconducting 
felt. The heat evolved by the change 
raises the temperature of the solution, 
the rise being indicated by the thermom- 
eters D, D. During the change the solution 
is stirred by the stirrer E. If the weight 
of the water is (say) 2570 g. and the rise 
in temperature is 1.5° centigrade, the heat 
evolved is 2570 x 1.5 = 3855 cal. 

Heat changes accompanying changes in the state of water. 
We have stated that in the conversion of a definite weight of 
ice into water, or water into steam, a definite amount of heat 
is required. Now that we have discussed the unit of heat these 
facts may be stated as follows : To convert 1 g. of ice at 0° 
into water at 0° requires 80 cal., tvhile to convert 1 g. of water 
at 100° into steam at 100° requires 537 cal. 

Chemical energy. A body may possess energy due to its 
motion or to its position. A piece of coal (or of any fuel) 
possesses energy due neither to its motion nor to its position 




\ .\AsW^'A^WA'AV..'AiAm'A'Ak'AW^ 



Fig. 37. A calorimeter, 
a form of apparatus 
used for determining 
the quantity of heat 
evolved in any definite 
chemical change 



58 



CHEMISTRY AND ITS USES 



but to its ability to undergo combustion. For in this process 
heat is evolved, and since heat is a form of energy, and energy 
cannot be created, it must be that in the process of combus- 
tion some other form of energy present in the coal and oxygen 
is converted into heat energy. This form of energy is called 
chemical energy. It is this form of energy in which the chemist 
is especially interested, for every chemical change is accom- 
panied by a change in chemical energy, and the energy changes 
are quite as important to the chemist as 
the changes in matter. Thus the value 
of any kind of coal depends primarily 
upon the chemical energy of the coal, 
for the greater the chemical energy, the 
greater the amount of heat that will be 
developed in the combustion of the coal, 
and it is the heat that we desire. Chemi- 
cal energy, then, may be denned as that 
form of energy which manifests itself when 
chemical action takes place. 

Ozone. It is sometimes possible to 
change the energy content of a sub- 
stance; and by so doing we change its properties. Thus, if 
oxygen is subjected to the action of an electric discharge a 
new gas, called ozone, is formed which has the characteristic 
odor noticed about electrical machines when in operation. 
Ozone is simply oxygen in which the normal energy content 
of the oxygen is greatly increased. It is not possible to con- 
vert more than a small percentage of oxygen into ozone, since 
the latter gas is unstable and changes back into oxygen. A 
point is soon reached, therefore, when the rate of formation 
of ozone is just equal to the rate of its decomposition. 

Ozone is a very strong oxidizing agent. Air containing 
ozone is sometimes used in certain manufacturing processes 
as well as in the purification of water. 




Fig. 38. Formation of 
ozone by the action of 
moist air on phosphorus 






MATTEK AND ENERGY 59 

Preparation of ozone. For commercial purposes ozone is prepared 
by the action of electric discharges through air. Its formation in 
the laboratory may be shown by partially covering with water a few 
pieces of stick phosphorus placed in the bottom of a jar (Fig. 38). 
The slow oxidation of the cold phosphorus is attended by the con- 
version of some oxygen into ozone. The presence of ozone in the air 
in the jar is soon indicated by its characteristic odor as well as by 
the property it possesses of imparting a blue color to strips of paper, 
A, previously dipped into a solution of potassium iodide and starch. 

Allotropic forms of matter. Just as oxygen and ozone are 
made up of the same kind of matter, so we shall find that many 
of the other elements occur in forms differing greatly in proper- 
ties, this difference being due to their different energy contents. 
Such forms of matter are known as allotropic forms. 

EXERCISES 

1. Is the fact that a candle disappears on burning at variance with 
the law of conservation of matter? 

2. The density of iron is 7.86. Compare the weight of any piece of 
iron with that of an equal volume of water. 

3. Name some common substances that cannot be melted without 
decomposition. 

4. Name some familiar substances that cannot be boiled without 
decomposition. 

5. Why does not all the water in a pond freeze in winter? 

6. Some people place tubs of water in their cellars in cold winters 
to prevent the freezing of vegetables stored in the cellars. Is such a 
procedure of any value ? 

7. Why does a block of ice melt so slowly even in warm air? 

8. Why does water sprinkled on the floor cool the room? 

9. Land near large volumes of water, such as the Great Lakes, is 
generally regarded as the best for growing fruits. Is there any reason 
for this belief ? 

10. At which place would most time be required to boil an egg hard — 
on the top of Pikes Peak or at Atlantic City ? 

11. When gases are compressed, heat is generated. Explain. 



60 CHEMISTRY AND ITS USES 

12. Name three crystalline and three amorphous substances. 

13. Give three illustrations of transformation of energy in your daily 
experience. 

14. (a) Give three ways of generating heat (b) What is the source 
of the heat energy in each case ? 

15. Gasoline used in motor cars contains what kind of energy? 

16. Name three kinds of energy purchased for use in our homes. 

17. How does the energy of a compound compare with the energy 
of the elements present in the compound? 

18. Can you suggest the source of the energy which you possess? 

19. Account for the fact that sometimes rain freezes almost instantly 
when it strikes the ground or any object and thus coats it with a layer 
of ice. 

20. How many calories of heat will be required to change 100 g. of 
ice at 0° into steam at 100°? 

21. 25 kg. of ice at 0° is placed in a refrigerator. Suppose that the 
ice melts and the temperature of the resulting water as it flows from 
the refrigerator is 10°. How many calories of heat have been absorbed ? 

22. In a certain experiment 2250 g. of water at 20° was contained 
in a calorimeter. After a reaction the temperature of the water was 
24°. How much heat was evolved in the reaction ? 

23. 1 kg. of water at 20° was cooled by adding 100 g. of ice (0°) and 
allowed to stand until the temperature became constant. Assuming 
that no heat is lost in radiation, calculate the approximate temperature 
of the cooled water. 



CHAPTER VII 

COMPOUNDS OF HYDROGEN AND OXYGEN: WATER AND 
HYDROGEN PEROXIDE 

Water 

Properties of water. Because of its great abundance and 
its importance to all forms of living organisms, a great deal of 
interest naturally attaches to hydrogen oxide, or water, as it is 
commonly called. Pure water is an odorless and tasteless liquid, 
colorless in thin layers, but having a bluish tint when observed 
through a considerable thickness. It solidifies at 0° and boils 
at 100° under the normal pressure of 1 atmosphere. When 
water is cooled, it steadily contracts until the temperature 
of 4° is reached ; at lower temperatures it expands. Water is 
remarkable for its ability to dissolve other substances and is 
the most general solvent known. Chemists usually employ 
aqueous solutions of substances rather than the substances 
themselves, since as a rule chemical action takes place more 
readily in solution. 

Occurrence of water. Water not only covers about five 
sevenths of the surface of the earth and is present in the 
atmosphere in the form of vapor, but it is also a common 
constituent of the soil, of many rocks, and of every form 
of animal and vegetable organism. Nearly 70 per cent of the 
human body is water. This is derived not only from the water 
which we drink but also from the food which we eat, most of 
which contains a large percentage of water. The table given in 
the chapter on foods (Chapter XXIX) shows the percentage 
of water present in some of the more common foods. 

61 



62 CHEMISTRY AND ITS USES 

Historical. Water was regarded as an element until 1781, 
when Cavendish showed that it is formed by the union of 
hydrogen and oxygen. Being a believer in the phlogiston 
theory, however, he failed to interpret his results correctly. 
A few years later Lavoisier repeated Cavendish's experiments 
and showed that water must be regarded as a compound of 
hydrogen and oxygen. 

Composition of natural waters. Water as it occurs in nature 
always contains more or less matter derived from the rocks 
and soils with which it comes in contact. When such water 
is evaporated this matter is left behind in solid form. Even 
rain water, which is the purest natural water, contains dust 
particles and gases dissolved from the atmosphere. The for- 
eign matter in natural waters is of two kinds ; namely, mineral 
and organic. 

1. Mineral matter. The mineral substances ordinarily pres- 
ent in fresh waters are common salt and compounds of calcium, 
magnesium, and iron. Water containing any considerable 
amounts of mineral matter does not form a lather with soap, 
and is termed hard water. Water containing little or no mineral 
matter, such as rain water, is termed soft water. One liter of an 
average river water contains about 0.175 g. of mineral matter. 
The waters of the ocean contain about 40 g. of mineral matter 
to the liter, more than three fourths of which is common salt. 

2. Organic matter. The organic matter present in water con- 
sists not only of the inanimate products of animal and vegeta- 
ble life but also of certain living microorganisms. This matter 
is absorbed from the soil or introduced from sewage. 

Effect of the foreign matter in water upon health. As a 
rule any sickness resulting from drinking impure waters is 
due to the presence of living microorganisms. Many of these 
are without injurious effect upon the human system, but some 
are the direct cause of disease. Thus a transmissible disease 
such as typhoid fever is due to a certain kind of organism 



WATER 



63 



which, through food or drink, is introduced into the system. 
It is easily possible for these organisms to find their way, through 
sewage, from persons afflicted with the disease into wells or any 
poorly protected water supply, and it is chiefly in this way that 
typhoid fever is spread. It may be added that the appearance 
of a water gives no conclusive evidence as to purity. A water 
may be unfit for drinking and yet be clear and odorless. 

The purification of water ; distilled water. Since all natural 
waters contain foreign matter, it is of interest to understand 
the methods used for 
removing such mat- 
ter and so obtaining 
chemically pure water. 
This is best done by 
the process of distilla- 
tion, which consists in 
boiling the water and 
then condensing the 
resulting steam. In 
the laboratory the 
process is usually con- 
ducted as follows : 




Fig. 39. The distillation of water as carried 
on in the laboratory- 



Ordinary water is poured into the flask A (Fig. 39) and boiled. 
The steam is conducted through the condenser B, commonly known 
as a Liebig condenser, which consists essentially of a narrow glass 
tube sealed within a larger one, the space between the two being 
filled with cold water, which enters at C and escapes at D. In this 
way the inner tube is kept cool and the steam in passing through 
it is liquefied. The water formed by the liquefaction of the steam 
collects in the receiver E and is known as distilled water. The 
impurities are not changed into vapor but remain in the flask A. 

Distilled water is pure water. It is used by the chemist in 
almost all his work. Large quantities are also used in the 
manufacture of ice, as well as for drinking. 



64 



CHEMISTRY AND ITS USES 



Commercial distillation. In preparing distilled water on a large 
scale the steam is generated in a metal boiler A (Fig. 40) and is 
conducted through the pipe B to the condensing coil C, made 
of tin. This pipe is wound into a spiral and is surrounded 
by cold water, which enters at D and flows out at E. The 
distilled water is collected in a suitable container F. 




Fig. 40. The distillation of water for commercial purposes 

Purification of water for sanitary purposes. When we 
speak of pure water, ordinarily we do not mean water from 
which all foreign matter has been removed but rather water 
which contains nothing injurious to health. All foreign matter, 
whether injurious or not, may be removed by distillation, but 
this process is costly, especially when we come to purifying 
large quantities of water. The method actually used for ren- 
dering water safe for drinking depends upon the volume of 
water to be purified. 

1. Purification of water on a small scale. If we have any reason 
to doubt the purity of a water used in our homes for drinking 



WATEE 



65 



\ ^ ^"^^ A 


, 












1^=^^- Impure water~-^=^^~_- 


g=gj 




"'..••.'_'• - . . Sand ■ 





we can easily render it safe by simply boiling it for a few 
minutes. The high temperature destroys any microorganisms 
present. These organisms, although destroyed by heat, can 
withstand very low temperatures without losing their vitality. 
2. Purification of water on a large scale. Many cities find it 
necessary to take their water supply from rivers and reser- 
voirs hi which the water is more or less contaminated with 
organic matter. Such a water supply is a source of constant 
menace to the 
health of a city. 
It becomes neces- 
sary, therefore, to 
find some way of 
effectively puri- 
fying the water 
on a large scale. 
Two general meth- 
ods are used for 
this purpose : (1) 
treatment with 
some germicide, 
usually chlorine 
(p. 141), which destroys the microorganisms present, and (2) fil- 
tration. As a rule the two methods are combined. The water 
is filtered on a large scale by allowing it to drain through 
large filtration beds made of sand and gravel. Some of the 
impurities are strained out by the filter, while others are de- 
composed by the action of certain kinds of microorganisms 
which collect in a jellylike layer on the surface of the filter. 
Such filters are known as slow sand filters. 

Fig. 41 shows a cross section of such a filter bed. The water 
filters through the sand and gravel and passes into the porous 
pipe A, from which it is pumped into the city mains. The niters 
are usually covered to prevent the water from freezing. 



Fig. 41. 



A covered slow sand filter for purifyin* 
water on a laro;e scale 



6Q 



CHEMISTRY AND ITS USES 



In some cases the water before nitration is run into large 
tanks and treated with certain compounds, such as alum, which 
form in the water a small amount of jellylike substances. This 
slowly settles to the bottom of the tank, carrying with it 
much of the organic matter present. The resulting water is 
then filtered through a so-called mechanical filter (Fig. 42) 
composed of sand and gravel. This method of treatment has 

largely replaced the 
slow sand filter, since 
the process is much 
more rapid and just 
as effective. 

The effect of the 
filtration of the water 
supply upon the health 
of a city is shown by 
the fact that typhoid 
fever is practically un- 
known in a city where 
all the water consumed 
is purified by an effect- 
Water taken from a well in 




Eig. 42. Mechanical filter in a modern city 
filtration plant 



ive purification plant (Fig. 43) 

a city is always a source of danger. 

Self-purification of water. It has long been known that water 
contaminated with organic matter tends to purify itself when 
exposed to the air (p. 26). This is due to the fact that air is 
somewhat soluble in water and that the dissolved oxygen, in the 
presence of certain microorganisms, gradually oxidizes the organic 
matter present in the water ; when this is destroyed, the organisms 
present die for lack of food. While water is undoubtedly puri- 
fied in this way, the process cannot be relied upon to purify a 
contaminated water so as to render it safe for drinking purposes. 

Chemical conduct. Water is a very stable substance ; in 
other words, it does not undergo decomposition readily. To 



WATER 



67 



decompose it into its elements by heat alone requires a very 
high temperature. Even at 2500° only about 10 per cent of 
the water heated is decomposed, and on cooling the constituent 
gases again combine to form water. Though very stable 
toward heat, water can be decomposed in other ways, as by 
the action of the 
electric current or 
by certain metals. 
Though contain- 
ing 88.81 per cent 



of 



oxygen, 



water 



is not a good oxi- 
dizing agent be- 
cause of its great 
stability. However, 
certain metals, as 
well as carbon, can 
be oxidized by very 
hot steam, the hy- 
drogenbeing set free 
in gaseous form. 



150 1904 


'05 '06 


'07 'C 


8 '09 


'10 


'11 


'12 


'13 


'14 


'15 


'16 


'17 


'18 


'19 


1920 


150 
































125 




























































































125 




























































100 
























































































































































75 




























































































75 




























































50 


- 


























































50- 
























































































25 


- 





























































































































































■ 


■ 
















0" 


" ■ 




1 


* 


t 


1 


T 












~m 




■ 






Unfiltered Water | Filtered Water 





Fig. 48. This figure gives the typhoid fever death 
rate per 100,000 of population in Columbus, Ohio, 
and shows the effect of water purification in de- 
creasing the ravages of this disease 



Water combines directly with many compounds, forming 
substances called hydrates. Blue vitriol and alum are good 
examples of such hydrates. A substance that contains no 
water whatever is said to be anhydrous. The term anhydrous 
literally signifies without water. 



Heat of formation and heat of decomposition are equal. The fact 
that a very high temperature is necessary to decompose water 
into hydrogen and oxygen is in accord with the fact that a great 
deal of heat is evolved by the union of hydrogen and oxygen 
(p. 35), for it has been proved that the heat necessary to decompose 
(say) 1 gram of a compound into its elements (heat of decomposi- 
tion) is equal to the heat evolved in the formation of 1 gram of the 
same compound from its elements (heat of formation). 



68 CHEMISTRY AND ITS USES 

Composition of water. We have already learned that water 
is composed of hydrogen and oxygen in the proportion of 
11.19 per cent of hydrogen to 88.81 per cent of oxygen 
(p. 18). The statement has also been made that the compo- 
sition of a compound can be determined only by experiment. 
The methods used for determining the composition of water 
will now be given, and the experiments will be described 
somewhat in detail, since they illustrate the methods used for 
determining the composition of compounds in general. 

Methods used in determining the composition of compounds : 
analysis and synthesis. In general we can use either of two 
methods in determining the composition of a compound: 
(1) after finding out what elements are present in the com- 
pound we may devise a way to decompose a definite weight 
of it into the elements which compose it, and accurately weigh 
the amounts of the elements so obtained ; or (2) we may 
measure the proportion in which the elements combine to form 
the compound. The former process is known as analysis, and 
this, in general, consists in separating a substance into its compo- 
nent parts ; the latter is known as synthesis, and this consists 
in the building up of a compound from its component parts. 

Each of the processes, analysis and synthesis, may be either 
qualitative or quantitative. Thus, qualitative analysis consists in 
separating a substance (either a compound or a mixture) into its 
component parts and finding out what these are. Quantitative 
analysis, on the other hand, consists in finding the weight of 
each component present in a definite weight of the substance. 

In determining the composition of water we might natu- 
rally think that a simple method would consist in decompos- 
ing water by an electric current and measuring the relative 
weights of oxygen and ' hydrogen evolved (p. 18). This 
method, however, is not very accurate. 

Proportion in which hydrogen and oxygen combine to form 
water. Experiments have already been described by which 






WATEK 



69 



water is proved to be a compound of hydrogen and oxygen 
(Figs. 19, 21). It remains for us to find out in what pro- 
portion these elements combine to form water. This may be 
done by mixing known volumes of hydrogen and oxygen, 
causing them to combine, and then ascertaining the volume 
and identity of the gas remaining. 

Details of the experiment. The combination of the two gases is 
brought about in a tube called a eudiometer. This is a graduated 
glass tube about 60 cm. long and 2 cm. wide, closed at one end 
(Fig. 44). Near the closed end two platinum 
wires are fused through the glass, the ends 
of the wires within the tube being sepa- 
rated by a space of 2 or 3 mm. The tube is 
entirely filled with mercury and inverted in 
a vessel of the same liquid. Pure hydrogen 
is passed into the tube until it is about one 
fourth filled. The tube is then lowered until 
the mercury stands at the same level inside 
and outside the tube, and the reading of the 
volume of the, hydrogen is taken. Approxi- 
mately an equal volume of pure oxygen is 
then introduced, and the volume is again 
taken. This gives the total volume of the two 
gases. From this the volume of the oxygen 
introduced may be determined by subtract- 
ing from it the volume of the hydrogen. 

The combination of the two gases is now 
brought about by connecting the two plati- 
num wires with an induction coil and passing a spark from one 
wire to the other. Immediately a slight explosion occurs. The 
mercury in the tube is at first depressed because of the expansion 
of the gases due to the heat generated, but it at once rebounds, 
taking the place of the gases which have combined to form water. 
The volume of the water in the liquid state is so small that it 
may be disregarded in the calculations. 

In order that the temperature of the excess gas and the mercury 
may become uniform, the apparatus is allowed to stand for a few 




Fig. 44. The eudi- 
ometer employed in 
determining the com- 
position of water 



70 CHEMISTKY AND ITS USES 

minutes and the volume of the gas is taken. The gas left over 
is then tested in order to ascertain whether it is hydrogen or oxy- 
gen, since experiments have proved that it is never a mixture of 
the two. From the information thus obtained the composition of 
the water may be calculated. 

Calculation of composition. In an experiment carried out with 
great exactness in the laboratory, the readings were as follows : 

Volume of hydrogen 20.3 cc. 

Volume of hydrogen and oxygen 38.7 cc. 

Volume of oxygen 18.4 cc. 

Volume of gas left after combination has taken 

place (found to be oxygen) 8.3 cc. 

We have thus found that 20.3 cc. of hydrogen have combined 
with 18.4 cc. minus 8.3 cc. (or 10.1 cc.) of oxygen, or approxi- 
mately 2 volumes of hydrogen have combined with 1 volume of 
oxygen. Since oxygen is 15.9 times as heavy as hydrogen, the 
proportion by weight in which the two gases combine is 1 part of 
hydrogen to 7.94 parts of oxygen. 

Morley's results. The American chemist Morley (Fig. 48) 
has determined the composition of water with greatest care. 
Extreme precautions were taken to use pure materials and to 
eliminate all sources of error. The hydrogen and oxygen which 
combined, as well as the water formed, were all accurately 
weighed. According to Morley's results 1 part by weight of 
hydrogen combines with 7.94 parts by weight of oxygen to 
form water. 

Comparison of results obtained. From the above discussion 
it is easy to see that it is by experiment alone that the com- 
position of a compound can be learned. Different methods 
may lead to slightly different results. The more accurate the 
method chosen, and the greater the skill with which the exper- 
iment is carried out, the more accurate will be the results. It 
is generally conceded by chemists that the results obtained by 
Morley in reference to the composition of water are the most 
accurate ones. In accordance with these results, then, water 



WATER 



71 



c 



WJ 



must be regarded, as a compound containing hydrogen and 
oxygen in the ratio of 1 part by weight of hydrogen to 7.94 
parts by weight of oxygen. 

Relation between the volume of water vapor and the volumes 
of hydrogen and oxygen which combine to form it. If the 
experiment described above (Fig. 44) is carried out at a tem- 
perature above 100°, the water vapor formed is not condensed, 
and it thus becomes possible to com- 
pare the volume of the water vapor 
with the volumes of hydrogen and 
oxygen which combined to form it. 
This can be accomplished by using a 
eudiometer of the shape of the letter 
U and surrounding the upper part of 
the eudiometer A (Fig. 45) with a 
glass tube B, through which is passed 
at C the vapor obtained by boiling 
some liquid which has a boiling point 
above 100°. This vapor keeps the 
tube A heated above the boiling point 
of water. In this way it has been 
proved that 2 volumes of hydrogen 
and 1 volume of oxygen combine to 
form exactly 2 volumes of water vapor. It tvill be noted that 
the relation between these volumes may be expressed by whole- 
numbers. The significance of this very important fact will be 
discussed in a subsequent chapter. 

Law of definite composition. We have just seen that water 
contains hydrogen and oxygen combined in a perfectly definite 
ratio. In the earlier days of chemistry there was much dis- 
cussion as to whether the composition of a given compound 
is always precisely the same or whether it is subject to some 
variation. Experiments have shown, however, that the com- 
position of a pure chemical compound is always exactly the 




Fig. 45. Eudiometer for 
measuring the volume of 
steam formed by the union 
of oxygen and hydrogen 



72 CHEMISTRY AND ITS USES 

same. Thus, pure water obtained from any source whatever, 
such as melting pure ice, condensing steam, or burning hydro- 
gen in oxygen, always contains 1 part by weight of hydrogen 
to 7.94 parts of oxygen. This truth is known as the law of 
definite composition and may be stated thus: The composition 
of a chemical compound never varies. 

Hydrogen Peroxide 

Composition. As has been shown, 1 part by weight of hydro- 
gen combines with 7.94 parts by weight of oxygen to form 
water. It is possible, however, to obtain a second compound 
of hydrogen and oxygen differing from water in composition in 
that 1 part by weight of hydrogen is combined with 2 x 7.94, 
or 15.88, parts of oxygen. This compound is called hydrogen 
peroxide, the prefix per- signifying that it contains more 
oxygen than hydrogen oxide (water). 

Properties and chemical conduct. Hydrogen peroxide is a 
clear, sirupy liquid whose freezing point is — 1.7°. It is diffi- 
cult to prepare in a pure state, since it is very unstable, decom- 
posing into water and oxygen with explosive violence : 

hydrogen peroxide >■ water -f- oxygen 

In dilute solution it is fairly stable, although it should be 
kept in a dark, cool place ; otherwise the solution loses its 
strength, the hydrogen peroxide present gradually decompos- 
ing into water and oxygen. The presence of a small percentage 
of certain substances, such as a trace of acid, preserves the 
strength by retarding decomposition. 

Uses. Solutions of hydrogen peroxide are used as oxidizing 
agents. The solution sold by druggists contains 97 per cent of 
water and 3 per cent of the peroxide. It has been largely used 
in medicine as an antiseptic, but recent experiments show it 
has little value for this purpose. It acts upon certain dyes 




Flg. 46. John Dal ton (English) (1766-1844) 

Teacher and scientist. He developed the atomic theory and made many- 
studies on the properties and composition of gases. His book, entitled 
"A New System of Chemical Philosophy," had a large influence on the 
development of chemistry 







Fig. 47. Theodore William 








Ry. "^W 


Richards (1868- ) 

Professor of chemistry in 
Harvard University and 




K£S# :4M£~49M 


well known for his exact 




E03&ti.' : *M*^ f "^IW 


investigations in many 




A # / 


branches of chemistry. 
The only American chemist 
to receive the Nobel prize 
($40,000) , which is bestowed 




V - ^lIHjiH 


annually upon those who 
have achieved greatest dis- 
tinction for their contribu- 
tions to knowledge 






Fig. 48. Edward Williams 




J^^ V 


Morley(1838- ) 




i*«4k 


Emeritus professor of 
chemistry in Western Re- 
serve University. Known 
for his accurate determina- 




WmLaf 


tions of the densities of 
oxygen and hydrogen and 
of the ratio in which they 




Sl 


combine to form water 




.■-.'. / 












HYDKOGEN PEROXIDE 73 

» 
and natural colors, such as that of the hair, oxidizing them 

to colorless compounds; hence it is sometimes used as a 

bleaching agent. 

Law of multiple proportion. It has been shown that both 

water and hydrogen peroxide are compounds of hydrogen 

and oxygen and that the ratio by weight in which these two 

elements are present in each of these compounds is as follows : 

Water hydrogen : oxygen = 1 : 7.94 

Hydrogen peroxide .... hydrogen : oxygen = 1 : 15.88 

It will be seen that the ratio between the weights of oxygen 
combined with a fixed weight of hydrogen (say 1 g.) in these 
two compounds is 7.94 : 15.88, or 1 : 2. 

Similarly, many elements other than oxygen and hydrogen 
unite to form a number of distinct compounds, each with its 
own precise composition. In all such compounds the same 
statement holds as in the case of water and hydrogen peroxide — 
the weights of the one element which are combined with a fixed 
weight of the other always bear a simple ratio to each other, 
such as 1 : 2 or 2 : 3. This truth is known as the law of multiple 
proportion. It was formulated by John Dalton (Fig. 46) in 1808 
and may be stated thus : When any two elements, A and B, 
combine to form more than one compound, the weights, of A ivhich 
unite with any fixed weight of B bear the ratio of small whole 
numbers to each other. 

EXERCISES 

1. What is the approximate weight of water in your body? 

2. What foods have a high water content (p. 268) ? What ones have 
a low water content ? 

3. In making solutions why does the chemist use distilled water 
rather than filtered water? 

4. How could you determine the total amount of solid matter dis- 
solved in a sample of water ? 

5. How could you determine whether a given sample of water is 
distilled water? 



74 CHEMISTRY AND ITS USES 

6. How could the presence of air dissolved in water be detected? 

7. How could the amount of water in a food such as bread or 
potato be determined? 

8. Would ice frozen from impure water necessarily be free from 
disease germs? 

9. Why is it that merely heating water to the boiling point is not 
sufficient to render it safe for sanitary purposes ? 

10. («) What are the advantages of using distilled water for drink- 
ing ? (b) Can you suggest any possible disadvantage ? 

11. Mention different ways in which a well water could become 
contaminated. 

12. What is your judgment concerning the probable purity of water 
taken from wells sunk in towns and cities ? 

13. If the water in your home town or city is purified, report on 
the method of purification. 

14. Why is cold water passed into C instead of D in Fig. 39 ? 

15. Mention at least two advantages that a metal condenser has 
over a glass condenser. 

16. Draw a diagram of the apparatus used in your laboratory for 
supplying distilled water. 

17. Suppose that the maximum density of water were at 0° instead 
of at 4°. What effect would this have on the formation of ice on bodies 
of water ? 

18. Compare the properties of water with the properties of the 
elements which compose it. 

19. State the two laws given in this chapter and illustrate each with 
an example. 

20. Distinguish between the terms analysis and synthesis; between 
hydrates and anhydrous substances. 

21. (a) Give the characteristics of an oxidizing agent, (h) What 
oxidizing agents have been mentioned so far in our study ? 

22. Give an example of a chemical action in which the speed of the 
action is increased by the presence of a small amount of some definite 
compound ; also one in which the speed is decreased. 

23. 20 cc. of hydrogen and 7 cc. of oxygen are placed in a eudiom- 
eter and the mixture is exploded. What gas, and how much of it, 
remains in excess? 



HYDROGEX PEROXIDE 75 

24. (a) What weight of oxygen is contained in 100 g. of water? 
(b) in 100 g. of pure hydrogen peroxide ? 

25. From the following data determine the ratio in which oxygen 
and hydrogen unite, the volumes all being measured under the same 
conditions of temperature and pressure : 

Volume of oxygen in eudiometer 8.54 cc. 

Volume of oxygen and hydrogen 52.72 cc. 

Volume of gas (hydrogen) left after explosion 27.10 cc. 

26. Morley found the composition of water by determining the 
weights of hydrogen and of oxygen that combine with each other to 
form water. The results of two trials are as follows : 

Hydrogen used Oxygen used 

(1) 3.2615 g. 25.9176 g. 

(2) 3.2559 g. 25.8531 g. 

In each case calculate the ratio in which the hydrogen and oxygen 
combined to form water. 

27. («) What weight of water would be formed by the combustion 
of 100 liters of hydrogen? (h) What volume of oxygen would be 
required in (a)? (c) What weight of potassium chlorate is necessary 
to prepare this amount of oxygen '? 



CHAPTER VIII 

NITROGEN AND THE RARE ELEMENTS ARGON, HELIUM, 
NEON, KRYPTON, XENON 

Properties of nitrogen. We have seen that oxygen is that 
constituent of the atmosphere which supports life. It is a very 
active element, however, and we can readily imagine what 
would happen if the atmosphere were all oxygen. This great 
activity of oxygen is kept in check by the presence in the 
atmosphere of a large percentage of the inert gas which we 
call nitrogen, which will neither burn nor support combus- 
tion. Like oxygen and hydrogen, nitrogen is a colorless, 
odorless, and tasteless gas. One liter of it weighs 1.25 g. It 
is very slightly soluble in water, 100 cc. of water dissolving 
but 2.33 cc. under standard conditions. It can be obtained in 
the form of a colorless liquid which boils at — 195.7°. 

Occurrence. Dry air is composed principally of oxygen and 
nitrogen in the free state, about 78 parts out of every 100 parts 
by volume being nitrogen. Nitrogen also occurs in nature 
combined with potassium and oxygen in the form of potassium 
nitrate (commonly called saltpeter or niter) ; it also occurs com- 
bined with sodium and oxygen in the form of sodium nitrate. 
Nitrogen is likewise an essential constituent of all living organ- 
isms ; it is therefore an essential constituent of our foods, being 
present in them in the form of protein (see chapter on Foods). 

Historical. Nitrogen was discovered by the Scottish chemist 
Rutherford in 1772. A little later the Swedish chemist Scheele 
(Fig. 53) showed it to be a constituent of air, and Lavoisier gave 
it the name azote, which means that it will not support life. The 
name nitrogen was afterwards given to it because of its presence " 
in niter. 

76 






NITROGEN AND THE RARE ELEMENTS 



77 



Preparation. Nitrogen can be readily obtained either from 
air or by decomposition of compounds containing the element. 

1. Preparation from air. On a large scale nitrogen is always 
prepared from liquid air, as will be explained in a subsequent 
chapter. In the laboratory it is prepared from air by the 
action of some substance which will combine with the oxygen, 
leaving the nitrogen free. Such a substance must be chosen, 
however, as will combine with the oxygen to form a product 
that is not a gas and that can be readily separated from the 
nitrogen. The substances most often used for this purpose 
are phosphorus and copper. 

(a) By the action of phosp)horus. 
The method used for the preparation 
of nitrogen by the use of phosphorus 
is as follows : 




Fig. 49. Preparing ni- 
trogen by burning out 
the oxygen of air with 
phosphorus 



The phosphorus is placed in a little 
porcelain dish supported on a cork and 
floated on water (Eig. 49). It is then 
ignited by contact with a hot wire, 
and immediately a bell jar or bottle is 
brought over it so as to confine a por- 
tion of the air. The phosphorus combines with the oxygen to 
form an oxide of phosphorus known as phosphorus pentoxide. 
This is a white solid which floats about in the bell jar, but which 
in a short time is all absorbed by the water, leaving the nitrogen. 
As fast as the oxygen is used up the water rises in the bell jar. 

(b) By the action of copper. The oxygen in the air may 
also be removed by passing air slowly through a heated tube 
containing copper. The copper combines with the oxygen to 
form copper oxide, which is a solid. The nitrogen passes on 
and may be collected over water. The details of the process 
are as follows : 

The copper is placed in a tube A (Fig. 50) and heated. Air is 
then forced slowly through the tube by pouring water into the 



78 



CHEMISTRY AND ITS USES 



bottle B. The oxygen of the air combines with the hot copper, 
forming the black solid, copper oxide, which remains in the 
tube, while the nitrogen passes on and is collected over water in 
the cylinder C. 

2. Preparation from compounds of nitrogen. In preparing 
nitrogen from air, the nitrogen is never entirely pure, since 
it contains the rare elements argon, helium, neon, krypton, and 
xenon, originally present in the air. These constitute only 
about 1 per cent of the nitrogen and, since they resemble the 




Fig. 50. Preparing nitrogen by removing the oxygen from air 
with hot copper 

nitrogen, do not materially affect its properties. Perfectly pure 
nitrogen is prepared by decomposing some of its compounds. 

Chemical conduct. Under ordinary conditions nitrogen is 
very inactive. It does not easily combine with oxygen, as is 
evident from the fact that the air contains both of these 
gases ; nor does it combine with anything else very readily, 
as is apparent from the fact that there is so much of it in 
the air and so little in the earth's crust in combination with 
other elements. 

At high temperatures the activity of nitrogen greatly in- 
creases. When it is mixed with oxygen and strongly heated 
a small fraction of the two gases combine, but the action is 



XITROGEX AXD THE EAEE ELEMENTS 



79 



always very incomplete. The best results are obtained by 
passing electric sparks through a mixture of the two gases 
(or air) or by causing the mixture to flow through an electric 
arc. Under these conditions nitric oxide (NO) forms. Under 
similar conditions nitrogen and hydrogen combine to a limited 
extent to form the compound known as ammonia. When 
nitrogen is heated with metals the action is much more ener- 
getic, particularly with magnesium, titanium, or aluminium. 
The resulting compounds are called nitrides, just as compounds 
of an element with oxygen are called oxides. 

Uses. Free nitrogen is used in the preparation of certain 
fertilizers. During the World War the Germans, being cut 
off from natural supplies of ni- 
trates, used nitrogen in the man- 
ufacture of ammonia, and this in 
turn in making nitric acid, which 
is essential for the manufacture of 
explosives. The gas is also used 
in certain gas-filled electric lamps. 

Assimilation of nitrogen by plants. 
While nitrogen is an essential con- 
stituent of both plants and animals, 
yet with the exception of a few plants none of these organ- 
isms have the power of directly assimilating free nitrogen. 
For example, the nitrogen in our bodies is taken from our 
foods and not from the air which we breathe. It has long been 
known, however, that certain plants, chiefly clover, alfalfa, 
beans, and similar plants belonging to the botanical order 
Leguminosoe, not only thrive in poor soil but at the same time 
enrich it. It is now known that these plants obtain at least a 
portion of their nitrogen from the atmosphere. This is accom- 
plished by groups of microorganisms which are gathered in 
little tubercles on the roots of the plants (Fig. 51). These 
organisms really constitute the workers in little chemical 




Fig. 51 . Tubercles on the roots 
of bean plants 



80 



CHEMISTRY AND ITS USES 



factories located on the roots of the plants and in which the 
nitrogen of the air is converted into compounds of nitrogen ; 
some of these compounds are assimilated by the plant and some 

are left in the soil and 
thus enrich it (Fig. 52). 

Argon, helium, neon, 
krypton, xenon. These 
are rare gaseous ele- 
ments and occur in the 
air in very small quan- 
tities. They are similar 
in that they are all col- 
orless, odorless gases. 
They differ from all 
other known elements 
in that they are entirely 
inactive, forming no com- 
pounds whatever. Argon, 
the most abundant of 
the group, was discov- 
ered in 1894 by two 
British scientists, Lord 
Eayleigh and Sir Wil- 
liam Eamsay (Fig. 54). 
It is now prepared from 
liquid air and, like nitrogen, used in certain gas-filled electric 
lamps. In 1868 Lockyer showed that a gaseous element, to which 
he gave the name helium, was present in the gases surrounding 
the sun. In 1895 Eamsay showed that this same element was 
present in the gases evolved in heating certain rare minerals and 
that traces were present in the atmosphere. Later Cady found it 
in the natural gas issuing from certain wells in Kansas and Texas. 
It is the most difficult of all gases to liquefy, having a boiling point 
of — 268.7°. The three remaining members of the group were dis- 
covered by Eamsay and Travers in 1898. They obtained them from 
liquid air, thus proving their presence in the atmosphere. Tubes 
containing neon are used to detect faulty automobile spark plugs. 




Fig. 52. A field of alfalfa 

One of nature's factories for the assimilation 
of nitrogen 



Fig. 53. Karl Wilhelm 
Scheele (1742-1786) 

This Swedish chemist was one 
of the greatest experimenters of 
his century. He discovered oxy- 
gen independently of Priestley, 
as well as the elements chlorine, 
tungsten, and molybdenum. He 
prepared the deadly gases arsine 
and prussic acid. He first iso- 
lated many organic compounds, 
such as glycerin and oxalic, tar- 
taric, and citric acids. Scheele 
was one of the ablest as well 
as one of the latest defenders 
of the phlogiston theory of 
combustion 




Fig. 54. Sir William Ramsay 
(1852-1916) 

An English chemist, who, to- 
gether with the English scientist 
LordRayleigh, discovered argon. 
In association with Travers he 
also discovered the elements 
xenon, neon, and krypton and 
was the first to show that helium 
exists on the earth. For these 
and other outstanding dis- 
coveries, Ramsay was awarded 
the Nobel prize in 1904 






t <& 


P 






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f 


v-- -ri 


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Fig. 55. An interior view of one of the plants built at Fort Worth, 
Texas, for the separation of helium from natural gas 

Until the year 1918 helium had been obtained in minute quantities only. 
When the United States entered the World War an effort was made to 
find some noncombustible gas for filling observation balloons, since the 
records showed that 95 per cent of the casualties resulting from the use 
of such balloons was due to the highly combustible character of the 
hydrogen with which they were filled. Helium was the only gas that 
at all met the requirements. An effort was made, therefore, to obtain it 
from the natural gas found in certain localities in Texas, which contained 
about 1 per cent of helium. To separate the helium, advantage was taken 
of its very low boiling point. The problem was exceedingly difficult. 
When the armistice was signed, however, 147,000 cu. ft. of nearly pure 
helium, stored in steel cylinders, was on its way to our armies, and the 
gas, which at the beginning of the war was almost a chemical curiosity 
and cost about $1700 per cubic foot, was being prepared in quantities at 
a cost of 10 cents per cubic foot. At the close of the war the government 
constructed at Fort Worth, Texas, a large plant for the separation of the 
helium from the natural gas of that locality. The first dirigible inflated 
with helium encircled the Capitol at Washington, December, 1921 



NITROGEN AND THE RARE ELEMENTS 81 

TABLE OF RARE ATMOSPHERIC ELEMENTS 





Helium 


Neon 


Argon 


Krypton 


Xenon 


Weight of 1 liter . . . 

Boiling point of liquid 
form 

Number of volumes in 
1,000,000 volumes of 
air (approximate) . . 


0.1782 g. 

- 268.7° 

4.00 


0.9002 g. 
-239° 

12.3 


1.7809 g. 
- 186° 

9400 


3.708 g. 
- 151.7° 

0.05 


5.851 g. 
- 109° 

0.006 



EXERCISES 

1. Contrast the properties and chemical conduct of oxygen, hydro- 
gen, and nitrogen. 

2. In what connections has Scheele's name been mentioned? 

3. Calculate the relative weights of nitrogen and oxygen; of nitro- 
gen and hydrogen. 

4. In preparing nitrogen from the air why not use a burning candle 
for removing the oxygen ? 

5. In Fig. 49 why does the water rise in the jar? Does this sug- 
gest a way for determining the percentage of oxygen in the air ? 

6. Compare the solubilities of oxygen, hydrogen, and nitrogen in 
water. 

7. Do plants really assimilate free nitrogen ? 

8. Why is alfalfa regarded as such an important crop in agriculture? 

9. State the significance of each of the following names : argon, 
helium, neon, krypton, xenon, nitrogen. (Consult dictionary.) 

10. In what respect do the rare elements in the atmosphere differ 
from nitrogen and all other elements? 

11. (a) 100 liters of dry air contains how many liters of nitrogen? 
(?>) What is the weight of this volume of nitrogen ? 

12. In an experiment 50 liters of dry air was passed over hot copper 
(Fig. 50). (a) What volume of nitrogen was obtained ? (&) How much 
did the copper increase in weight? 

13. A student prepared 25 liters of nitrogen collected over water in 
a laboratory in which the thermometer registered 20° and the barom- 
eter 742 mm. Calculate the volume of this nitrogen under standard 
conditions. 



CHAPTER IX 
MOLECULAR WEIGHTS ; ATOMIC WEIGHTS 

Molecules. We have seen that the way gases act, as 
described in the gas laws, proves beyond any doubt that 
every gas is made up of extremely small particles called 
molecules. Methods have been found for actually calculating 
the weights of the individual molecules with considerable 
accuracy, but these weights are so small that we can make 
little practical use of them. To help us realize how small 
the molecules are Langmuir has calculated that if the mole- 
cules in one cubic inch of air were to be each one changed 
into a grain of fine sea sand, the resulting sand would fill a 
trench a mile wide and three feet deep, reaching from New 
York to San Francisco. 

Atoms. A moment's reflection will show that the molecule 
of a compound, such as water, must be made up of different 
parts, for a compound contains at least two different ele- 
ments, and so each molecule of the compound must have 
some of each constituent. These constituent parts of a molecule 
are called atoms. The law of definite composition states the 
fact that the composition of any pure sample of a compound 
is always the same, so the composition of each molecule of 
the compound must be the same, since a weighable sample of 
the compound is merely a vast collection of molecules. It is 
easy to see that the weight of each atom of any one kind in 
all the molecules must be the same as that of any other, or 
the individual molecules would have different compositions, 
and we would not have the law of definite composition. 

82 



MOLECULAR WEIGHTS; ATOMIC WEIGHTS 83 

The atomic theory. It was this striking law of definite compo- 
sition that led John Dalton (Fig. 46) in 1806 to advance the theory 
that each elementary body is made up of a vast number of indi- 
vidual atoms, and this suggestion was called the atomic theory. 
We now have such convincing evidence that this picture of the 
make-up of matter is a true one that we accept it as a truth rather 
than as a theory, for any theory becomes a truth when there is no 
longer any doubt as to its reality. 

Relative molecular weights. It will be recalled (p. 47) that 
the simplicity of the gas laws led Avogadro to believe that 
equal volumes of any two gases inclose the same number of 
molecules, no matter what their size or weight may be. So 
many additional facts are now known that there can be no 
doubt that Avogadro's principle is a statement of the truth. 
We can even determine the number of molecules in a cubic 
centimeter of air with more precision than we can take the 
census of a large city. 

With the help of Avogadro's principle we can at once deter- 
mine the relative weights of the molecules of various gaseous 
substances. For in any definite volume — say 1 liter — there 
will always be the same number of molecules no matter what 
gas fills the liter measure. So the weights of the different 
liters will be in the same ratio as those of the several kinds 
of molecules constituting the gases. 

Thus a liter of oxygen weighs 1.429 g., a liter of hydrogen 
weighs 0.08987 g., and one of nitrogen weighs 1.2507 g. These 
figures give the actual weight in grams of equal numbers of the 
three kinds of molecules. They must therefore be in the same 
ratio to each other as the actual weights of the three kinds 
of molecules. 

Relative atomic weights. The next question is, Have we 
any simple way for determining the relative weights of the atoms 
that constitute the molecules ? At first sight it might appear 
that the law of definite composition affords us such a method. 



84 CHEMISTRY AND ITS USES 

Taking water as an example, we know that each portion we 
analyze consists of 1 part of hydrogen to 7.94 parts of oxygen 
by loeight. Therefore the composition of each molecule must be 
expressed by this ratio. If the molecules were all made up of 
one atom each of oxygen and hydrogen, then these two num- 
bers, 1 and 7.94, would be the relative weights of the atoms, 
and our problem would be solved. These numbers that state 
the ratio by iveight in which two elements combine are called the 
combining weights of the elements, and they can be determined 
with great precision by the analysis of compounds. 

But there is another compound of hydrogen with oxygen, 
namely, hydrogen peroxide, and in this the ratio is 1 : 15.88 
(p. 73), so that oxygen and hydrogen have two combining 
weights. Here our trouble begins, for we have no direct way 
to tell which of these two compounds (if either of them) is 
the one that has one atom of each element in its molecules. 
It is evident, however, that in all such cases the one combining 
ratio (or combining iveight) must be a simple multiple of the other, 
for the atoms will have to combine in some definite small ratio, 
such as 1 : 1 or 1 : 2. This fact at once shows the reason for 
the law of multiple proportion. 

We have now made real progress in determining the rela- 
tive weights of the atoms, for we can determine the ratio by 
weight of two elements in the compounds they form with each 
other (the combining weights), and some one of these combining 
iveights must be the atomic weight. Years of effort to solve the 
problem completely have shown that we must follow an unex- 
pected plan. We must first get the molecular weights of com- 
pounds of the elements and from these draw conclusions as 
to the relative weights of the constituent atoms. 

Standard for molecular weights. When we have to deal with 
any series of purely relative values, such as molecular weights, 
we usually choose some one as unity and express the others 
in terms of this one. For example, in our coinage we may select 



i 



MOLECULAR WEIGHTS; ATOMIC WEIGHTS 85 

the dollar as our standard and state* the value of the other 
coins as fractions and multiples of this ; or we may select the 
penny and state the value of all others as multiples of this. 
We may select the dime and call it 10 or 25 or any number 
we choose, and calculate the value of all other coins to the 
standard we have selected. Whatever coin we select and 
whatever value we give it, the ratio in value between all the 
coins will remain unchanged. 

In a similar way we might call oxygen unity. The ratio by 
weight between the liter of oxygen, hydrogen, and nitrogen 
would then be 1 : 0.0629 : 0.875. But in general we should 
like to avoid numbers that are less than unity, just as we 
have avoided coining a halfpenny, so we might take the liter 
of the smallest weight (hydrogen) as unity, and the ratios 
would then be 15.88 : 1 : 13.9. Since oxygen plays such an 
important part in chemistry we should like to stick to it as our 
standard, while keeping the weights of all other molecules 
above unity. We can therefore provisionally adopt the round 
number 16 for ox} r gen, and the ratio becomes 16 : 1.008 : 14.01. 

By determining the weight of one liter of each of all known 
gases we can decide upon the relative weights of their mole- 
cules referred to this standard (oxygen = 16) or any standard 
we may later adopt. 

Standard for atomic weights. Since the molecules are made 
up of atoms, and we are going to find a way to determine the 
relative weights of the atoms as well as those of the molecules, 
it is important to adopt a final standard for molecular weights 
that will be suitable for both molecules and atoms. Thus, if 
it were to turn out that the molecules of oxygen gas and 
hydrogen gas are not single atoms, but are each made up of tivo 
atoms, then the standard of 16 for the oxygen molecule would 
give us 8 for the oxygen atom and 0.5 for the hydrogen atom, 
and it would be better to adopt 32 for the oxygen molecule so 
as to have' 'the hydrogen atom at least 1. 



86 CHEMISTRY AND ITS USES 

Two atoms in the oxygen molecule. It is clear that if mole- 
cules of compounds consist of two or more atoms, the molecule 
of elementary gases, such as oxygen or hydrogen, may either 
be single atoms or they may be composed of two or more 
atoms of the same kind. That the molecules of both oxygen 
and hydrogen consist of two atoms can be shown as follows: 

When oxygen and hydrogen combine to form steam we 
have the volume relations shown in the equation (p. 71) 

2 vol. hydrogen + 1 vol. oxygen >- 2 vol. steam 

Let us suppose that the 1 volume of oxygen contains 100 mole- 
cules. Then the 200 volumes of steam must contain 200 mole- 
cules (Avogadro's principle). But each of these 200 molecules 
must contain at least one atom of oxygen, or 200 in all, and 
these 200 atoms came from 100 molecules of oxygen. There- 
fore each molecule of oxygen must contain at least two 
atoms. Similar reasoning shows that the molecule of hydro- 
gen also must contain at least two atoms. 

Evidently we have merely shown that there are at least 
two atoms in these molecules. There might be more than 
that, but there is no evidence that points this way, and we 
assume that there are two only. To retain the value 16 
for the atom of oxygen (and a little above unity for that of 
hydrogen) we will now adopt oxygen as 32 as the final 
standard for molecular and atomic weights and compare all 
others with this value. 

Molecular weights from the weight of 22.4 liters. Having 

adopted the arbitrary value oxygen = 32 as our standard, let 

us calculate the volume occupied by 32 g. of oxygen gas. 

Since 1 liter weighs 1.429 g. the volume occupied by 32 g. 

32 
will be M Ann i or 22.4 liters. If we make a vessel that holds 
1.429 

exactly 22.4 liters and fill it with oxygen, it will hold the stand- 
ard molecular weight (measured in grams); namery, 32 g. If 



MOLECULAE WEIGHTS; ATOMIC WEIGHTS 87 



we now replace the oxygen with nitrogen, the same number of 
molecules will be present as before (Avogadro's principle), and 
experiment shows that the weight is 28.02. But these two 
numbers, 32 and 28.02, are the relative weights of an equal 
number of molecules, so that the nitrogen molecule weighs 
28.02 compared with the oxygen molecule as 32. In like man- 
ner the weight of 22.4 liters of any gas will give its molecular 



224 
LITERS 



/ 


/ 


22.4 
LITERS 





y<. 


y 


22.4 

LITERS 


y 



S s 


22.4 
LITERS 


^ 



oxygen, 32 g. hydrogen, 2.016 g. nitrogen, 28.016 g. water vapor, 18.016 g. 

Fig. 56. The weight of 22.4 1. of various gases 

weight referred to oxygen as 32 (Fig. 56), and we can say 
that the number which expresses the weight of 22.4 liters of any 
gas is the same as the number which expresses the molecular 
weight of the gas. 

Remarks. Very few of the gases expand and contract exactly 
as the gas laws say they should. Consequently the weight of 
22.4 liters of a gas is not its exact molecular weight, but is 
within a few tenths of one per cent of it. 

Evidently the method we have found applies merely to gases 
(including liquids and solids that can be converted into gases). 
In an elementary course we cannot take up any other cases, 
and we need only say that methods are known for determining 
the molecular weights of liquids and of dissolved solids. 

Selection of the atomic weight from the combining weights. 
It will now be easy to show how we can decide which one of 
the combining weights of an element is really its atomic weight, 
and the best way to do this is by an example. So let us 
suppose that we have found that the combining weights of 
nitrogen are 7.005, 14.01, and 21.015; the atomic weight is, 
therefore, one of these numbers. 



88 



CHEMISTRY AND ITS USES 



We first determine the weight of 22.4 liters of a number of 
gaseous compounds which we know to contain nitrogen. These 
values are given in the first column of the following table : 



Gaseous Compounds 


Molecular 

Weight 
(22.4 Liters) 


Percentage of 
Nitrogen by 
Experiment 


Part of Molecu- 
lar Weight due 
to Nitrogen 


Nitrogen gas . . . . 
Nitrous oxide .... 
Nitric oxide .... 
Ammonia ...... 


27.95 
44.13 
40.00 
17.05 


100.00 
63.70 
46.74 

82.28 


27.95 
28.11 
14.02 
14.03 



We next analyze each of these compounds to ascertain the 
percentage of nitrogen present, placing the values obtained in 
the second column. If we multiply the molecular weight of 
each compound by the percentage of nitrogen, the product will 
be the portion of the molecular weight due to nitrogen. But since 
the molecules are made up of atoms, the part of a molecule 
due to nitrogen must represent the sum of the weights of the 
nitrogen atoms present. We notice that the numbers in the 
last column are either very near to 14 or to twice 14 and that 
none are near 7 or 21. If we examine a large number of nitro- 
gen compounds, it is reasonable to expect that we should find 
some that contain only one atom, and since we find none which 
give a value of less than 14, we assume that this and not 7 
or 21 or 28 represents the weight of a nitrogen atom. 

Accurate determination of atomic weights. The weight of 
a given volume of a gas is difficult to determine with great 
precision, and in consequence the molecular weights of gases 
as determined by experiment are usually subject to a very 
considerable error. The portion of nitrogen in 22.4 liters of 
the various gases is therefore a little uncertain, as will be seen 
from the values in the table above. All these figures tell us 
is that the true value is very near 14. The combining weight 
can be very accurately determined by the analysis of any of 



MOLECULAR WEIGHTS; ATOMIC WEIGHTS 89 

these compounds, and is found to be 7.005. It is therefore 
evident that the accurate atomic weight is twice this value; 
namely, 14.01. 

Summary. These, then, are the steps which must be taken 
to establish the atomic weight of an element : 

1. Determine the combining weight accurately by analysis. 

2. Determine the weight of 22.4 liters of a large number of 
gaseous compounds of the element and, by analysis, the part 
of the molecular weights due to the element. Tlie smallest 
number so obtained will be the approximate atomic weight. 

3. Multiply the combining weight by the integer (1, 2, or 3), 
which will give a number close to the approximate atomic 
weight. The number so obtained will be the precise atomic ic eight. 

Molecular weights of the elements. When we determine 
the weight of 22.4 liters of the various elementary gases we 
reach some interesting conclusions. Experiment shows that 
the molecular weights of many of them, such as nitrogen, 
chlorine, and bromine, give values which are twice the atomic 
weights, so that -in these cases the molecule contains two atoms, 
as we have found to be true with oxygen and hydrogen (p. 86). 
In the case of the metals, so far as their vapors have been 
studied, the molecular weight and the atomic weight are the 
same, so that the molecide of a meted must consist of a single 
atom. The molecule of ozone, on the other hand, contains 
three atoms of oxygen, while the molecules of phosphorus 
and arsenic contain four atoms. 

Weight of a liter of gas. The weight of 1 liter of each gas is 
determined by actually weighing the gas, and from this weight 
we calculate the molecular weight of the gas. Conversely, if 
we happen to know the molecular weight of the gas, we can 
easily calculate back to the weight of 1 liter of the gas, for the 
number which represents the molecular weight also represents 
the weight of 22.4 liters of the gas. Hence, to find the weight 
of 1 liter of any gas divide its molecular weight by 22.4. 



90 CHEMISTRY AND ITS USES 

EXERCISES 

1. Is it correct to speak («) of an atom of a compound? (b) of a 
molecule of an element ? 

2. State the law upon which is based our method for determining 
molecular weights. 

3. Almost every one of us uses the word theory. What do we mean 
by the term ? 

4. Is a theory an expression of truth ? If not, what is the distinction? 

5. Mention a theory that has been discarded because proved false 
(p. 20), also one that has become a truth. 

6. (a) Mention the reasons why oxygen is selected as a standard 
for atomic and molecular weights rather than hydrogen, which is the 
lightest of all elements, (b) Why not represent oxygen as 1 rather than 
16 as the standard for atomic weights ? 

7. Calculate the relative weights of the molecules (a) of hydrogen 
and oxygen (see Appendix for weights of 1 liter of various gases) ; 
(6) of hydrogen and nitrogen. 

8. Calculate the molecular weights of hydrogen, oxygen, and hydro- 
gen chloride from the weight of 1 liter of each of these gases (Appendix). 

9. Carbon dioxide is one of the gases exhaled from our lungs. Cal- 
culate its molecular weight from the weight of 1 liter of the gas. 

10. When sulfur burns in oxygen or air (p. 20) a gas known as sul- 
fur dioxide forms. The molecular weight of sulfur dioxide is 64.06. 
(a) Calculate the weight of 1 liter of the gas. (b) Compare your result 
with that given in the Appendix. 

11. The chief constituent of natural gas is a compound of carbon 
and hydrogen known as methane. Its molecular weight is 16.04. (a) Cal- 
culate the weight of 1 liter of the gas. (b) Compare your result with 
that given in the Appendix. 

12. Sulfur dioxide contains 50.05 per cent of sulfur. One liter of 
the gas weighs 2.9266 g. Assuming that the molecule of the sulfur 
dioxide contains but one atom of sulfur, calculate the atomic weight of 
sulfur from the above data. (Suggestion. First calculate the molecular 
weight of the gas ; then determine how much of this weight is sulfur.) 

13. Carbon dioxide (see exercise 9, above) contains 72.72 per cent 
of oxygen. Assuming that the molecule of the gas contains two atoms 
of oxygen, calculate the atomic weight of oxygen. 






CHAPTER X 
FORMULAS; EQUATIONS; SOLUTION OF PROBLEMS 

Percentage composition. Just as we can determine the com- 
position of water with great accuracy (p. 68), so, by similar 
means, we can determine the composition of other compounds. 
Having analyzed a given compound, we usually express its com- 
position in percentages, or in the parts of each element in 100 
parts of the compound. Thus, we have seen that water consists 
of 88.81 per cent of oxygen and 11.19 per cent of hydrogen. 
This mode of expression takes no account of the fact that com- 
pounds are made up of molecules, the atoms of which each 
have characteristic weights. It would be much better to have a 
method of stating composition which will express all these facts. 

Atomic composition. Remembering that the atomic weight 
of oxygen is 16, it is evident that if we divide the percentage 
of oxygen in water by 16, the quotient (5.55) will be the 
relative number of oxygen atoms in 100 parts of water. In like 
manner, if we divide the percentage of hydrogen (11.19) by 
the atomic weight of the element (1.008), the quotient (11.10) 
will express the relative number of hydrogen atoms in 100 parts 
of iv at er. The two numbers, 5.55 and 11.10, therefore repre- 
sent the ratio between the number of oxygen and hydrogen atoms 
in 100 g. of water. But this same ratio must hold for any other 
quantity of water, even for one' molecule, since any quantity of 
water is made up of molecules. To reduce the ratio to its sim- 
plest terms we divide the two numbers by the smaller one : 

5.55 - 5.55 = 1 ; 11.10 - 5.55 = 2 

The ratio of oxygen atoms to hydrogen atoms in a molecule 
of water is therefore 1 : 2. 

91 



92 CHEMISTRY AND ITS USES 

Formulas. We may express the ratio found for the oxygen 
and hydrogen atoms in a molecule of water by writing the 
two symbols together thus, H 2 0, the subscript 2 indicating 
that two atoms of hydrogen are in combination with one of 
oxygen. The ratio of oxygen atoms to hydrogen atoms 
(namely, 1 : 2) could, however, be just as well represented 
by any multiple of H 2 0, such as H 4 2 or H 6 3 , for in all 
these the ratio of the oxygen to hydrogen atoms is the same. 
We must therefore decide, if possible, which of these ex- 
presses the true composition of the molecule of water. Now 
it is very easy to do this if we know the molecular weight of 
water, for if H 2 is right, then the molecular weight of water 
must be equal to 2 times the atomic weight of hydrogen 
(2 x 1.008 = 2.016) + the atomic weight of oxygen (16), or 
18.016. If H 4 2 is correct, then the molecular weight will be 
2 x 18.016, or 36.032. Now, it is possible by Avogadro's prin- 
ciple to determine the molecular weight of water vapor, and 
in this way it is found to be approximately 18. Hence H 2 
correctly represents the composition of the molecule of water 
as vapor. The expression H 2 is termed the formula of water. 
It may then be stated in general that the formula of a com- 
pound or element is an expression which represents the composi- 
tion of the molecule of the compound or element in terms of atoms. 

Derivation of the formula of potassium chlorate. Let us take 
another example. Potassium chlorate when analyzed in the labo- 
ratory is found to be as follows : potassium, 31.9 per cent ; chlorine, 
28.9 per cent ; oxygen, 39.2 per cent. Now, proceeding as in the 
case of water, we get the following results : 

31.9 -s- 39.10 = 0.8158 = relative number of atoms of K in 100 g. 
28.9 + 35.46 = 0.8150 = relative number of atoms of CI in 100 g. 
39.2 -f- 16.00 = 2.4500 = relative number of atoms of O in 100 g. 

Dividing the three quotients by the smallest (0.8150) we get the 
integers 1, 1, 3. (Since analyses are always slightly inaccurate, 
the ratios found will often differ slightly from whole numbers, but 



FORMULAS 93 

the difference is so slight as to leave no doubt as to what the inte- 
ger really is.) The simplest formula of potassium chlorate is 
therefore KC10 3 , and its real formula is either KC10 3 or some 
multiple of this. If KC10 3 is correct, then its molecular weight is 
(39.10 -f- 35.46 + 3 x 16), or 122.56. Experiments show that the 
molecular weight is approximately 122 ; hence KC10 3 is the correct 
formula. 

Calculation of the molecular weight and the percentage 
composition of compounds. We have seen that it is possible to 
determine the percentage composition of a compound and its 
molecular weight by experiment, and knowing these we can cal- 
culate the formula of the compound. On the other hand, hav- 
ing once determined the formula of a compound and knowing the 
atomic weights of the elements present in it (see Appendix), we 
can easily calculate back to the molecular weight and the per- 
centage composition from which the formula was derived ; and 
it is sometimes convenient to do this, for it is easier to remember 
formulas than percentage composition. For example, knowing 
that the formula of water is H 2 0, its molecular weight must 
be the sum of two atomic weights of hydrogen (2 x 1.008) and 
one.atomic weight of oxygen (16), or 18.016. Now if the mole- 
cule of water weighs 18.016 and contains one atom of oxygen 
weighing 16, then 16/18.016, or 88.81 per cent, of the water is 
oxygen, and 2.016/18.016, or 11.19 per cent, must be hydrogen. 

Let us take another example ; namely, hydrogen sulfate. Its 
formula is H 2 S0 4 ; hence its molecular weight is the sum of 2 times 
the atomic weight of hydrogen (2 x 1.008) plus the atomic weight 
of sulfur (32.06) 4- 4 times the atomic weight of oxygen (4 x 16), 
or 98.076. 

Percentage of hydrogen = ' ■ = 2.05 

32 06 

Percentage of sulfur = ' = 32.70 

64 

Percentage of oxygen = = 65.25 



94 CHEMISTRY AND ITS USES 

Gram-molecular weights. For practical purposes we deal 
with pounds or with grams of a substance, not with atoms 
and molecules. Now, since the numbers 18.016, 16, and 2.016 
represent the ratio by weight between a molecule of water and 
the oxygen and hydrogen of which it is composed, the same 
ratios must hold between any weight of water we may choose 
and the oxygen and hydrogen in this weight of water. Evi- 
dently in 18.016 lb. of water there will be 16 lb. of oxygen and 
2.016 lb. of hydrogen, and in 18.016 g. there will be 16 g. of 
oxygen and 2.016 g. of hydrogen. 

For practical purposes, therefore, we may allow the symbol H 
to stand for 1.008 g. of hydrogen, the symbol O for 16 g. 
of oxygen, and the formula H 2 for 18.016 g. of water. 
The weight in grams of an element, corresponding to its 
atomic weight, is called a gram-atomic weight. The weight 
in grams of an element or of a compound, corresponding to 
its molecular weight, is called a gram-molecular iveight. 

Equations. Having devised a convenient way of expressing 
the composition of compounds, not in percentages but in for- 
mulas, we make use of equations to express chemical trans- 
formations, using an arrow in place of an equality sign. For 
example, the equation 

2H + 0— ^H 2 (1) 

is a concise method of stating two distinct facts. 

1. Qualitatively, it states that water is formed by the union 
of hydrogen and oxygen. 

2. Quantitatively, it tells us that 2 gram-atomic weights of 
hydrogen (2.016 g.) combine with 1 gram-atomic weight of 
oxygen (16 g.) to form 1 gram-molecular weight of water 
(18.016 g.). 

Molecular equations. Since a formula expresses the com- 
position of a molecule, and since experiment has shown that a 
molecule of oxygen and one of hydrogen each contain two 



EQUATIONS 95 

atoms, the formulas of these gases are written 2 and H 2 rather 
than 2 O or 2 H, which would simply represent two atoms not 
combined. If we wish our equation to state these additional 
facts, we shall have to change it to the form 

2H 2 + 2 ^2H 2 (2) 

This is called a molecular equation, and it will be seen that 
it expresses the same ratios by weight as does equation (1). 
It also expresses the fact that 2 molecules of hydrogen com- 
bine with 1 molecule of oxygen to form 2 molecules of water, 
and this for some purposes makes it a more useful equation. 
In this text we shall ordinarily use the molecular equations, 
although we may sometimes use the simple form if we wish 
only to express ratios. 

While the molecules of hydrogen and oxygen are made 
up of two atoms each, this is not the case with all elements. 
Thus the phosphorus molecule contains four atoms, and we 
write it P 4 . In the case of some elements, especially the metals, 
the molecule contains but one atom ; in such cases, therefore, the 
atom and the molecule are identical. 

Decomposition of potassium chlorate. Let us take another example 
of the meaning of an equation. It will be remembered that oxygen 
was prepared by heating potassium chlorate, which has the formula 
KC10 3 . When heated, this compound decomposes into oxygen and 
a compound called potassium chloride whose formula is KC1. The 
decomposition is represented by the equation 

2 KC10 3 ►■ 2 KC1 + 3 2 

This equation states the following facts : 

1. Qualitatively, potassium chlorate decomposes into potassium 
chloride and oxygen. 

2. Quantitatively, 2 gram-molecular weights of potassium chlo- 
rate (2 x 122.6 g.) decompose into 2 gram-molecular weights of 
potassium chloride (2 x 74.56 g.) and 3 gram-molecular weights of 
oxygen (3 x 32 g.). The number before a formula applies to the 



96 CHEMISTRY AND ITS USES 

formula as a whole, while the subscript number applies only to 
the symbol which it follows. 

3. Molecularly, 2 molecules of potassium chlorate decompose 
into 2 molecules of potassium chloride and 3 of oxygen. 

Steps in writing an equation. We must keep in mind that 
chemical equations simply express facts deduced from experiments. 
In other words, we cannot find out by equation-writing what 
the formula of a compound is or what changes the compound 
undergoes, but having found out these facts we can express 
them by an equation. In writing an equation the first step 
consists in writing the formulas (or symbols) of the substances 

entering into the reaction on the left of the arrow ( >-), 

and on the right the formulas (or symbols) of the substances 
formed. Thus, having learned that the formula of mercuric 
oxide is HgO, that it decomposes on heating into mercury and 
oxygen, and that the oxygen molecule is 2 , while the mercury 
atoms remain single, we express these facts as follows : 

First step : HgO >■ Hg + 2 

The second and final step consists in balancing the equation ; 
that is, we must modify our equation (if necessary) so that 
there will be just as many atoms of each element on one side 
of the equation as on the other. Thus, in the above equation 
we have two atoms of oxygen (or 1 molecule) on one side of the 
equation. To get this we must take 2 HgO, and this will give 
us 2 Hg as well as 2 ; hence we write the completed equation, 
keeping in mind also that the molecules and atoms of metals 
are as a rule identical (p. 95) : 

Second step : 2 HgO >- 2 Hg + 2 

Equations of reactions so far studied. Let us now put into 
the form of equations the reactions studied up to this point, 
writing both steps. Remember that the complete equation in each 
case is the correct one ; the first is written simply as an aid in 
formulating the completed equation. 



EQUATIONS 97 



1. Preparation of oxygen : 

(a) From potassium chlorate: 
First step, 

KCIO3 — *KC1 + 2 
Complete, 

2KC10 3 ^2KCl + 30 2 

(5) From the electrolysis of water : 

First step, 

H 2 0— >H 2 + 2 
Complete, 

2H 2 0— v2H 2 + 2 

2. Preparation of hydrogen : 

(a) From sodium and water: 
First step, 

Na + H 2 >- NaOH + H 2 

Complete, 

2 Na + 2 H 2 >■ 2 NaOH + H 2 

(6) From zinc and sulfuric acid: 

First step and complete (identical), 

Zn + H 2 S0 4 ^ZnS0 4 + H 2 

(<?) From zinc and hydrochloric acid: 
First step, 

Zn + HCl— ^ZnCl 2 + H 2 
Complete, 

Zn + 2 HC1 y ZnCl 2 4- H 2 

(rf) From iron and steam : 
First step, 

Fe + H 2 0— >-Fe s 4 + H 2 
Complete, 

3 Fe + 4 H 2 — >- Fe 8 4 + 4 H 2 



98 CHEMISTRY AND ITS USES 

Representation of the heat of reaction. We can also employ 
chemical equations to express the heat given off or absorbed 
during chemical action. The equation 

2 H 2 + 2 — > 2 H 2 + 136,800 cal. 

states the fact that when 4.032 g. of hydrogen combines with 
32 g. of oxygen, forming 36.032 g. of water, heat is given off 
to the extent of 136,800 cal. Evidently when 1 gram-molecular 
weight (18.016 g.) of water is formed, 68,400 cal. is given 
off, and this is called the heat of formation of water. 

Conditions of a reaction not indicated by equations. Equa- 
tions merely state the composition of the substances taking part 
in the reaction and the weights of each one involved, together 
with the energy change measured as heat. They do not tell 
the conditions under which the reaction will take place. For 
example, the equation 

2HgO— v2Hg + 2 

does not tell us that it is necessary to keep heating the mer- 
curic oxide at a moderately high temperature in order to effect 
its decomposition. The equation 

Zn + H 2 S0 4 ►■ ZnS0 4 + H 2 

in no way indicates that the hydrogen sulfate must be dissolved 
in water before it will act upon zinc. The equation 

s + o 2 — ^so 2 

does not indicate that no perceptible action takes place unless 
the sulfur is first heated, but that when once started it goes 
on of its own accord and with a bright flame. 

It will therefore be necessary to pay close attention to the 
details of the conditions under which a given reaction occurs, 
as well as to the statement of the equation itself. 

Problems based on equations. Since an equation is a state- 
ment of the weights of materials which take part in a reaction, 



SOLUTION OF PROBLEMS 99 

when the equation has once been established by experiment 
we can use it in calculating the various weights. A few 
examples will show how this may be done. 

1. How many grams of oxygen are evolved on heating 100 g. 
of mercuric oxide ? 

First write the equation for the reaction involved : 

2HgO ^2Hg + 2 (1) 

Next determine the relative weights of the amounts of the differ- 
ent substances involved in the reaction. The atomic weights of 
mercury and oxygen are respectively 200.6 and 16 (see table 
on back cover). Hence the relative weight of the 2 HgO equals 
2(200.6 + 16), or 433.2. Similarly, the relative weight of the 
oxygen evolved, namely, 2 , equals 2 x 16, or 32. It is convenient 
now to write these numbers under the formulas in equation (1). 

This then becomes n TT _ ^ _ TT _ 
2HgO )-2Hg + 2 

433.2 32 

These numbers indicate that 433.2 units by weight (in this 
case grams) of mercuric oxide will, on heating, evolve 32 units by 

32 

weight of oxygen ; hence 1 g. of mercuric oxide will give g. 

on 4:66.2 

of oxygen, and 100 g. will give 100 x , QQO » or 7.38 g. 

4:66.2 

2. I wish to prepare 100 g. of oxygen, using potassium chlorate 
as a source of the oxygen. How many grams of the chlorate will 
be required ? 2 KCl09 ^ 2 KC1 + 3 2 

245.12 96 

The figures tell us that 245.12 g. of potassium chlorate will 

yield 96 g. of oxygen ; hence to prepare 1 g. of oxygen we must 

245.12 
use — — — g. of potassium chlorate, and to prepare 100 g. we 

245 12 
must have x 100 = 255.33 g. 

yo 

3. How many grams of zinc must be dissolved in sulfuric acid 
to produce 10 g. of hydrogen ? 

Zn + H 2 S0 4 ^ZnS0 4 + H 2 

65.37 ' 2.016 



100 CHEMISTRY AND ITS USES 

To obtain 2.016 g. of hydrogen we must use 65.37 g. of zinc. 

Hence to obtain 1 g. of hydrogen, * g. of zinc is required, 

and to obtain 10 g. we must take 10 times as much ; namely, 
ggx 10 = 324.25. 

It must be remembered that the equations show relations by 
weight , not by volume ; hence in problems involving volumes of 
gases it will be necessary first to find the weights of the gases. 
The table in the Appendix gives the weight of 1 liter of each of 
the common gases, measured under standard conditions. The 
following problem will illustrate the method: 

4. How many grams of potassium chlorate are necessary to 
prepare 100 liters of oxygen? 

Since 1 liter of oxygen weighs 1.429 g., 100 liters will weigh 
142.9 g. 2 KC10 ^ _^ 2 KC1 + 3 2 

245.12 96 

^p x 142.9 = 364.86 g. 

It will be recalled that in Chapters III and IV we have 
solved a number of problems of the same kind as those solved 
above, but in the earlier chapters these problems were solved 
by referring to the percentage composition of the compounds. 
Having now learned the significance of formulas and equations, 
it is much simpler to solve problems of this kind by reference to 
the formulas and equations involved in the problem. 

In working out such problems do not carry divisions beyond the 
second decimal place. Having completed a problem, look to see 
if the result is reasonable. 

EXERCISES 

1. State all the facts expressed by the formula H 2 ; by H 2 S0 4 . 

2. State all the facts expressed by the equation 

2H 2 ^2H 2 + 2 

3. A compound analyzed in the laboratory was found to contain 
2..76 per cent hydrogen and 97.24 per cent chlorine. Its molecular 
weight was found to be approximately 36. Calculate its formula. 



SOLUTION OF PROBLEMS 101 

4. Common salt has a molecular weight of approximately 58 and 
contains 39.34 per cent sodium and 60.65 per cent chlorine. Calculate 
its formula. 

5. Ordinary saltpeter has a molecular weight of approximately 101 
and contains 38.67 per cent potassium, 13.88 per cent nitrogen, and 
47.45 per cent oxygen. Calculate its formula. 

6. A compound has the following composition : hydrogen, 5.92 per 
cent ; oxygen, 94.07 per cent. Its molecular weight is approximately 34. 
What is the compound? 

7. (a) Hydrogen sulfate has the formula H 2 S0 4 . Calculate its 
molecular weight and its percentage composition from its formula. 
(b) Compare your results with those given on page 33. 

8. (a) Do both of the following equations balance? 

(1) KC10 3 >-KCl + 30 

(2) 2 KC10 3 >■ 2 KC1 + 3 2 

(b) What additional facts does the second one express ? 

9. Balance the following equations : 

Ca + HCl ^CaCl 2 + H 2 

Fe + HCl ^FeCl 2 + H 

10. Suppose you wish to prepare 80 g. of oxygen by heating 
potassium chlorate : («) write the equation involved in the reaction ; 
(b) from the equation calculate the weight of potassium chlorate that 
will be required. 

11. A student added to 100 g. of zinc sufficient sulfuric acid to dis- 
solve it. (a) Write the equation for the reaction involved ; (b) calculate 
the weight of hydrogen evolved ; (c) calculate the volume of the 
hydrogen. 

12. When iron is dissolved in hydrochloric acid, hydrogen is evolved, 
as expressed in the equation 

Fe + 2 HC1 >■ FeCl 2 + H 2 

Suppose you wish to prepare 100 liters of hydrogen by this method. 
(a) What would this volume of hydrogen weigh ? (b) What weight of 
iron would be necessary to prepare this weight of hydrogen ? 

13. A student heated 25 g. of potassium chlorate in a flask until all 
the oxygen was evolved, (a) Calculate the volume of the oxygen given 
off. (b) What weight of potassium chloride remained in the flask? 



CHAPTER XI 
CARBON AND ITS OXIDES 

Introduction. The term carbon is more or less familiar to 
everyone. It suggests coal and smoke and troubles with auto- 
mobile engines. In the cities, especially, it annoys us by soiling 
our clothes. Not so many persons, however, realize that this 
soft black substance which is so useful as a fuel has also its 
artistic value, for it is the same substance which composes 
the very hard, brilliant gem prized for centuries and known 
as the diamond. Carbon also occurs in other forms, and before 
studying its properties we must first learn what some of these 
forms are. 

Occurrence. In the uncombined state carbon is found in 
nature in several forms. The diamond is practically pure 
carbon, while graphite and the various forms of coal all con- 
tain more or less free carbon. The element also occurs abun- 
dantly in the form of compounds. Carbon dioxide, which we 
breathe from our lungs, is its most familiar gaseous com- 
pound. Natural gas and petroleum are largely compounds 
of carbon and hydrogen. It is a constituent of limestone, 
as well as of many other rocks. All living organisms, both 
plant and animal (p. 10), contain a large percentage of this 
element, and the number of its compounds which go to make 
up the vast variety of animate nature is almost limitless. In 
the free state carbon occurs in both the crystalline and the 
amorphous form. 

Crystalline carbon. Crystalline carbon occurs in two forms ; 
namely, diamond and graphite. jV -**-" 

"~I^"i02 



CAKBON AND ITS OXIDES 



103 




Fig. 57. The largest diamond ever 

found, the famous Cullinan diamond, 

in its original condition (one half 

natural size) 



1. Diamond. Diamonds are found in certain localities in 
South Africa, the East Indies, and Brazil. The crystals as 

found are usually covered 
with a rough coating. These 
are cut so as to bring out the 
brilliancy of the gem. Dia- 
mond-cutting is carried on 
very extensively in Holland. 

Fig. 57 is a photograph of 
the largest diamond ever found, 
in its original condition. It 
is known as the Cullinan dia- 
mond and was presented to 
King Edward VII by the 
Transvaal government. Fig. 58 

is a photograph (natural size) of another very famous diamond 

(the Kohinoor), in finished form. 

The density of the diamond is 3.5, and, though brittle, it is 
one of the hardest of substances. Few chemical reagents have 
any action on it; but when heated 
in oxygen or the air, it blackens 
and burns, forming carbon dioxide. 

Preparation of artificial diamonds. 
Many attempts have been made to 
produce diamonds artificially. For a 
long time these ended in failure, 
graphite and not diamonds being the 
product obtained, but in 1893 the 
French chemist Moissan (Fig. 112), in 
his study of chemistry at high tem- 
peratures, finally succeeded in making 
some small ones. He accomplished this by dissolving carbon in 
melted iron and plunging the crucible containing the mixture 
into water, as shown in Fig. 59. Under these conditions the 
carbon crystallized in the iron in the form of the diamond. The 




Fig. 58. The Kohinoor dia- 
mond after being cut (natural 
size) 



101 



CHEMISTRY AND ITS USES 



crystals were then freed from the metal by dissolving away the 
iron in hydrochloric acid. The resulting diamonds were too 

small to be of any use. 

2. Graphite. This form of carbon 
is found in large quantities, es- 
pecially in Ceylon, Siberia, and in 
some parts of the United States, 
Mexico, and Canada. Large quan- 
tities are also made commercially 
by heating hard coal to a high 
temperature. It is a shining black 
substance, very soft, and greasy 
to the touch. Its density is about 
2.15. It is used in the manufac- 
ture of lead pencils and crucibles, 

as a lubricant, and (in the form of a polish or a paint) as a 

protective covering for iron. 

Commercial production of graphite. The process consists in heating 
hard coal in large electric furnaces about 40 ft. in length, a longi- 
tudinal section of one of which is shown in Fig. 60. The elec- 
trodes A are made of graphite. The furnace is nearly filled with 




Fig. 59. Sketch illustrating 

Moissan's method of producing 

diamonds in the laboratory 




Fig. 60. Electric furnace for the production of graphite 

the coarse grains of coal B. Since the coal is a poor conductor of 
electricity, there is placed in the center of the charge a core, C, 
of carbon, which serves to carry the current through the charge. 
The charge is covered with a mixture, D, of sand and carbon 



CAEBON AND ITS OXIDES 



105 



(or similar materials), which serves to exclude the air. An electric 
current is supplied by the generator G. The current in passing 
through the charge meets with resistance sufficient to produce a 
high temperature. Under the influence of the intense heat the 
carbon is changed into graphite. Prepared in this way, the 
product is uniform in composition and free from grit and is 
therefore superior to the natural product for most purposes. 




Press Illustrating Service, Inc. 

Fig. 61. Mining coal in a typical mine 

Amorphous carbon. Pure amorphous carbon is best prepared 
by heating sugar (CjH O u ) in the absence of air. The hydro- 
gen and oxygen present are expelled, largely in the form of 
water, and pure carbon remains. Almost any form of organic 
matter yields carbon when heated in the absence of air. Among 
the numerous substances that contain amorphous carbon, the 
following may be mentioned : 

1. Coal and coke. The various forms of coal (Fig. 61) were 
formed from vast accumulations of vegetable matter. In hard 
coal, or anthracite, nearly all the carbon present is in the uncom- 
bined state ; while in soft, or bituminous, coal a considerable 



106 



CHEMISTRY AND ITS USES 



portion of the carbon present is combined with hydrogen, 
oxygen, nitrogen, and sulfur. When soft coal is heated in the 
absence of air complex changes occur (as will be described in 
a later chapter), resulting in the formation of various useful 
compounds of carbon, which are given off in the form of gases 
and vapors, while the mineral matter and free carbon remain 
and constitute ordinary coke. The matter which escapes when 




Fig. 62. 



Drawing of a plant for the production of charcoal by the 
modern method 



coal is heated in the absence of air is known as volatile matter. 
In hard coal the volatile matter averages from 5 per cent to 
8 per cent, while in soft coal it averages from 30 per cent to 
35 per cent. When coal is burned, the mineral matter present 
is left in the form of ash. 

2. Charcoal. This is prepared from wood just as coke is 
prepared from coal. The volatile matter formed and expelled 
by the heat contains many valuable substances, such as wood 
alcohol and acetic acid, wh'ich are obtained commercially in this 
way. Formerly much of this volatile matter was allowed to 
escape, but at present a large amount of charcoal is prepared 






CARBON AND ITS OXIDES 107 

in such a way that the volatile matter is condensed and saved, 
as in the heating of coal for making coke. Both charcoal and 
coke are used as fuels and are especially useful in reducing 
metals from their oxides, as will be explained later. 

Modern methods for the production of charcoal. Iron cars are 
loaded with wood A, A (Fig. 62) and run into a low, narrow iron 
room, or retort B. The retort is then made air-tight and heated 
slowly for twenty-four hours by the tires F, F. The volatile prod- 
ucts escape through the pipes C, C and then pass into the con- 
densers D, D. Here those portions which are liquid at ordinary 
temperatures, such as wood alcohol and acetic acid, are condensed, 
while the gaseous products are led back and burned in the fires F, F. 
When all the volatile matter has been expelled in this way, the cars 
containing the charcoal are run into cooling chambers, and their 
place in the retort is taken by other cars loaded with wood. 

3. Boneblack, or animal charcoal. This is made by charring 
bones and animal refuse. It consists of very finely divided 
carbon and of calcium phosphate and is especially useful for 
removing coloring matter in the refining of sugar. 

4. Carbonblack ; lampblack. The black powders known as 
carbonblack and lampblack are products of the imperfect 
combustion of carbonaceous fuels, such as oil and gas. 

Destructive distillation. The process of heating a substance 
such as coal, wood, and bones, in the absence of air, is known 
as destructive distillation. Thus we say that charcoal is pre- 
pared by the destructive distillation of wood. This process 
differs from ordinary distillation in that complex changes 
occur, resulting in the formation of new compounds. 

Properties of carbon. While the various forms of carbon 
differ in many properties, especially in hardness, yet they are 
all odorless, tasteless solids, insoluble in water, and character- 
ized by their stability toward heat. Even when subjected to 
the intense heat of the electric arc, carbon does not. melt, 
though it vaporizes to some extent. 



108 



CHEMISTKY AND ITS USES 



In the form of boneblack carbon withdraws certain kinds of 
organic matter from solutions ; hence its use in refining sugar 
and other substances. A marked property of carbon is its ability 
to take up (or, in chemical terms, to adsorb') large volumes of 
gases. The amount of gas adsorbed by any given sample of car- 
bon depends not only on the 
material from which the car- 
bon is prepared but also on the 
method of preparation. That 
prepared from coconut shells 
and peach seeds is one of the 
best for this purpose. Thus, 
one volume of charcoal pre- 
pared from coconut shells will 
adsorb about 250 volumes of 
air and about 335 of chlo- 
rine, a gaseous element used 
so largely as a poison gas in 
the World War. Charcoal pre- 
pared especially for adsorbing 
gases is known as activated 
charcoal. 




Fig. 63. A form of gas mask used 

in the World War as a protection 

against poison gas 



Gas masks. The use of various 
poison gases such as phosgene 
(p. 26) in the World War made it necessary to devise gas masks 
for the protection of the troops. Many different kind of masks were 
used, but as finally developed, these were made and fitted to the 
face (Fig. 63) in such a way that all the inhaled air had to pass 
through a light metal box (canister) filled with layers of various 
materials, the chief of which was activated charcoal. These mate- 
rials either adsorbed or combined with the poison gases. They 
were so efficient that poison gas present in the ratio of 1000 parts 
of gas to 1,000,000 of air could be reduced to one part per million in 
the small fraction of a second required for the air to pass through 
the canister. 



CARBON AND ITS OXIDES 109 

Chemical conduct. At ordinary temperatures carbon is a 
very inert substance, but at higher temperatures it combines 
directly with most of the elements. Because of its strong 
affinity for oxygen it is an excellent reducing agent. Its com- 
pounds with the metals are called carbides. One of the most 
important of these is calcium carbide (CaC 2 ), which is used in 
the -preparation of acetylene. When carbon or a substance con- 
taining it, such as wood or coal, burns, the element combines 
with oxygen to form either carbon dioxide (C0 2 ) or carbon 
monoxide (CO). Both of these oxides are colorless gases. 

Uses of carbon. The chief use of amorphous carbon is for 
fuel to furnish heat and power for all the uses of civilization. 
An enormous quantity of carbon in the form of coal, coke, 
and charcoal is used as a reducing agent in the separation of 
the various metals from their ores. Carbonblack is used for 
making motor-car tires, indelible ink, printer's ink, paints, and 
black varnishes, while boneblack and charcoal are used in 
niters. In the refining of sugar the dark solution of the impure 
compound is filtered through boneblack, which removes the 
coloring matter. On evaporation the resulting solution yields 
the colorless sugar. Activated charcoal is used in refining 
various products and in making gas masks for protecting work- 
men in certain industries from the evil effects of poisonous gas 
evolved in industrial operations. # 

How the chemist determines the percentage of carbon in a sample of 
coal. To determine the percentage of carbon in a sample of coal 
the chemist weighs out a small amount of the sample, say 1 g., in 
a narrow dish A (Fig. 64) and slips it into the glass tube B, 
arranged so that it can be heated in the furnace, as shown in 
the figure. The remaining part of the tube is filled with some 
oxidizing agent, usually copper oxide. The glass tube is then 
gradually heated to a red heat, while a slow current of air, dried 
and free from carbon dioxide, is passed into the tube at C. Under 
these conditions the carbon in the coal burns to carbon dioxide. 
This gas then passes through the narrow tube D into the tube E, 



110 



CHEMISTRY AND ITS USES 



filled with calcium chloride, which absorbs any water present. 
The carbon dioxide passes on into F and there bubbles through 
a solution of some substance, such as sodium hydroxide, and is 




Fig. 64. Apparatus for determining the percentage of carbon in coal 

absorbed. The increase in the weight of the tube F (and con- 
tents) gives the weight of the carbon dioxide formed. From this 
weight of carbon dioxide the weight of carbon present in it is 
easily calculated, and this equals the weight of carbon present 
in the 1 g. of coal burned. 

Carbon dioxide (carbonic acid gas) : properties and occur- 
rence. Carbon dioxide, often known as carbonic acid gas, is 
a colorless and practically odorless gas which is exhaled 

from our lungs and is 

whenever any 

common fuels 

which contain 

burn in oxy- 




formed 
of the 
(all of 
carbon) 
gen or air 

C + o 



CO, 



Fig. 65. A method for showing that carbon 
dioxide is a heavy gas 



Large quantities also 
escape from volcanoes 
and crevices in the earth. 
It is present in the open 
air to the extent of about 3 parts in 10,000. When we study 
the atmosphere we shall see that this relatively small amount 
is essential to the life of all plants. 



CARBON AND ITS OXIDES 



111 



Carbon dioxide is one of the heaviest of gases, 1 liter weighing 
1.9768 g. Its weight may be inferred from the fact that it can 
be poured like water from one vessel downward into another ; 
or, as a more striking experiment, the gas may be poured from 
a cylinder A into a beaker B attached to a balance and counter- 
poised as shown in Fig. 65. At 15° and under ordinary pres- 
sure 1 volume of water dissolves 1 volume of the gas. It is 
rather easily condensed to a colorless solid, which, under ordi- 
nary pressure, evaporates without melting at — 78.5°. The gas 
is not regarded as poisonous, 
although life is not possible 
in an atmosphere containing 
large percentages of it. 

Liquid and solid carbon dioxide. 

The commercial carbon dioxide 

compressed in steel cylinders is 

under such great pressure that 

it is largely in the liquid state. 

When the pressure is removed 

the rapid vaporization of the 

liquid reduces the temperature 

sufficiently to freeze a portion of the escaping liquid to a snowlike 

solid (Fig. 66). Cylinders of liquid carbon dioxide are inexpensive 

and should be available in every school. It is possible to separate 

the carbon dioxide formed in the combustion of coal from the 

other gases formed at the same time, and this constitutes the chief 

commercial source of the gas. 

To prepare the solid carbon dioxide the cylinder should be 
placed across the table and supported in such a way that the stop- 
cock end is several inches lower than the other end. A loose bag 
is made by holding the corners of a piece of cloth around the neck 
of the stopcock. The stopcock is then turned on so that the liquid 
rushes out in large quantities. A considerable quantity of the snow 
very quickly collects in the cloth. 

Preparation. In the laboratory carbon dioxide is prepared 
by the action of hydrochloric acid on the compound known as 




Fig. 66. Carbon dioxide in the 
solid form 



112 



CHEMISTRY AND ITS USES 



calcium carbonate (CaC0 3 ). The latter is found in nature in 
many different substances, such as shells, coral, and limestone. 
Marble is nearly pure calcium carbonate and is the material 
most often used in the preparation of carbon dioxide. When 
hydrochloric acid and marble are brought in contact with each 
other, water, calcium chloride (CaCl 2 ), and carbon dioxide are 
formed according to the following equation : 

CaC0 3 + 2 HC1 — y CaCl 2 + H 2 + C0 2 

With sulfuric acid the products are calcium sulfate (CaS0 4 ), 
water, and carbon dioxide. 

In the preparation of carbon dioxide pieces of marble are placed 
in the flask A (Fig. 67). Hydrochloric acid is added drop by 

drop through the funnel tube B. The 
carbon dioxide escapes through C and, 
being heavier than air, collects in the 
cylinder, as shown in the figure. The 
calcium chloride formed is a white 
solid which remains dissolved in the 
water in A. 




Fig. 67. A simple apparatus 

for preparing carbon dioxide 

by the action of hydrochloric 

acid on marble 



Chemical conduct. Carbon dioxide 
is a very stable substance. It is 
neither combustible nor a supporter 
of combustion. When it is passed 
into a clear solution of calcium 
hydroxide, Ca0 2 H 2 (ordinary lime- 
water), the solution soon becomes 
cloudy, owing to the formation of 
calcium carbonate. The calcium carbonate (CaC0 3 ) is insol- 
uble and separates as a solid as fast as it is formed, producing 
a cloudy or milky appearance in the solution : 

Ca0 2 H 2 + C0 2 y CaC0 8 + H 2 

These properties constitute a simple test for carbon dioxide. 



CARBON AND ITS OXIDES 



113 



When passed into water a portion of the gas unites with 
the water, forming a very unstable compound (H 2 C0 3 ) known 
as carbonic acid. Thus : 



H.O+CO, 



H,CO ( 



Uses. The carbon dioxide in the air is a food for plants, as 
will be shown in the chapter on the atmosphere. Commercially 
it is used chiefly in the manufac- 
ture of soda water and similar bev- 
erages and as a fire extinguisher. 
Ordinary soda water consists of 
various flavoring extracts to which 
is added water charged with car- 
bon dioxide under pressure. When 
the pressure is removed, the excess 
of gas escapes, producing efferves- 
cence. A burning candle is extin- 
guished in air which contains as 
little as 2.5 per cent of carbon 
dioxide, hence the dioxide is used 
as a fire extinguisher. 




Fig. 68. A modern fire extin- 
guisher in which water and 
carbon dioxide are used to ex- 
tinguish the flame 



Portable fire extinguishers. Famil- 
iar types of portable fire extinguish- 
ers are shown in Figs. 68, 69. That 
shown in Fig. 68 is a device for 

generating carbon dioxide under pressure. The liquid is a solution 
of sodium bicarbonate in water contained in a metal cylinder. The 
bottle A contains sulfuric acid. In ease of fire the apparatus is 
grasped by the handle D, the apparatus is inverted, and the knob 
B is pushed in by tapping it against the floor. This breaks the 
bottle containing the sulfuric acid, which at once reacts with the 
sodium bicarbonate, generating carbon dioxide. Some of the gas 
dissolves in the water, while the rest forces the water out through 
the nozzle C. While the volume of water so obtained is not large, 
it is very effective as a fire extinguisher because of the carbon 



114 



CHEMISTKY AND ITS USES 



dioxide accompanying it. The type of extinguisher shown in 
Fig. 69 is a smaller metal cylinder filled with the liquid known 
as carbon tetrachloride (CC1 4 ) and fitted with a sort of piston 
for forcing the liquid out in case of fire. The liquid is not com- 
bustible and is quite volatile ; hence when it is thrown onto a fire 
it volatilizes and, the vapor being heavy, sur- 
rounds the burning body and prevents the oxy- 
gen of the air from coining in contact with it. 

Carbon monoxide (CO). This compound 
resembles carbon dioxide in that it is a 
colorless, nearly odorless gas. It differs 
from it in being very inflammable as well 
as very poisonous. 

It is interesting to note that birds are very 
sensitive to this gas. In mine explosions car- 
bon monoxide is always formed, and rescuers 
often carry canaries with them, the death of the 
birds warning the rescuers of their own peril. 

Carbon monoxide gas is formed whenever 
carbon is burned in a limited supply of 
oxygen: 2 C + 0„— v2C0 








Fig. 69. A fire extin- 
guisher for spraying 
carbon tetrachloride 
on a flame 



It is often formed in stoves when the air 
draft is shut off, especially when hard coal is 
used as a fuel. Since the gas is very poisonous, 
care should be taken that the pipes and chimneys are not closed in 
any way ; otherwise the gas may escape into the room and cause the 
death of those inhaling it. The gas is formed in the combustion of 
gasoline in engines, especially if the carburetor is not well adjusted, 
and a number of cases of poisoning have resulted from breathing 
the air in closed garages in which an engine is allowed to run. 

In the laboratory the gas is usually prepared by heating 
formic acid (H 2 C0 2 ) or oxalic acid (H 2 C 2 4 ) with sulfuric acid, 
which is added to absorb the water formed : 



HCO, 



H 2 



CO 



CARBON AND ITS OXIDES 115 

Chemical conduct. Chemically carbon monoxide is quite 
active. It combines readily with oxygen and burns in air with 
a characteristic pale-blue flame (often observed hi the com- 
bustion of hard coal), forming carbon dioxide: 

2CO + 2 — ^2C0 2 

It reduces metallic oxides such as copper oxide (CuO), forming 
the metal and carbon dioxide 

CO + CuO ^C0 2 + Cu 

Because of this property carbon monoxide is often used as a re- 
ducing agent in liberating the metals from their oxides. It com- 
bines with chlorine to ioim phosgene (COC1 ), a very poisonous, 
colorless gas used largely hi the World War as a poison gas. 
Solution of problems. We have seen that a gram-molecular 
weight of any pure gas is the weight of 22.4 liters of the gas 
expressed in grams. By keeping this in mind it is possible 
for us to calculate in a very simple way the volume relations 
of gases entering into a reaction. All that we have to do is to 
write the equation so that each gas is represented hi a molecu- 
lar state. The number of gram-molecules of each gas entering 
into the reaction when multiplied by 22.4 gives the volume of 
the gas. The following problem will make the method clear : 

What volume of oxygen is required to burn 100 liters of carbon 
monoxide ? First write the simple equation for the reaction : 

CO + o — y C0 2 

This, however, represents but 1 atom of oxygen ; hence it must 
be doubled so as to represent a molecule of oxygen (0 2 ) : 

2 CO + 2 y 2 C0 2 

This equation tells us that 2 gram-molecules of carbon monoxide, 
or 44.8 liters (2 x 22.4), will combine with 1 gram-molecule of oxy- 
gen, or 22.4 liters. In other words, 2 liters of carbon monoxide 
will combine with 1 liter of- oxygen ; hence 100 liters of carbon 
monoxide will combine with 50 liters of oxygen. 



116 CHEMISTRY AND ITS USES 

EXERCISES 

1. Suggest a method for proving that all the various forms of carbon 
described are really carbon. 

2. How could you determine the relative percentages of carbon in 
different samples of coal ? 

3. To what is the black color of the letters on this page due? 

4. Why does sugar turn black when heated? 

5. Why is the end of a telegraph pole often painted with carbon 
paint before placing it in the ground ? 

6. How does carbon differ from the other elements so far studied? 

7. Is there any relation between the value of different samples of 
coal and the ash which they form when burned ? 

8. How does boneblack differ in composition from carbonblack ? 

9. Would charcoal used in a gas mask ever lose its effectiveness? 

10. Contrast the properties of the two oxides of carbon. 

11. Why is carbon monoxide regarded as a treacherous poison? 

12. How could you distinguish between oxygen, hydrogen, carbon 
dioxide, and carbon monoxide ? 

13. When carbon monoxide acts as a reducing agent what change 
does it undergo ? 

14. Why are different methods used for collecting oxygen and carbon 
dioxide when they are prepared in the laboratory ? 

15. Carbon dioxide sometimes collects in deep wells. How could you 
prove its presence in such cases ? 

16. What are the sources of the carbon and the oxygen in the carbon 
dioxide which we breathe from our lungs ? 

17. Carbon dioxide contains a large percentage of oxygen, but it is 
not an oxidizing agent. Explain. 

18. Wood alcohol and acetic acid (as well as many other compounds) 
are obtained when certain kinds of wood are heated in the absence of air. 
Does this imply that these compounds are present in wood ? 

19. How do we account for the widely different properties of the 
diamond and coal? 

20. W^hat is the percentage of carbon (a) in carbon dioxide? "(b) in 
carbon monoxide? 



CAKBON AND ITS OXIDES 117 

21. («) Calculate the weight of oxygen combined with 1 g. of carbon 
in carbon dioxide. (Suggestion. First write the equation for the forma- 
tion of carbon dioxide and then proceed as explained in the paragraph 
entitled " Problems based on Equations," p. 98). (b) Calculate the 
weight of oxygen combined with 1 g. of carbon in carbon monoxide. 
Are the results obtained in (a) and (b) in accord with the law of multiple 
proportion ? 

22. The average person exhales about 8 liters of air per minute con- 
taining about 4 per cent of carbon dioxide by volume. Calculate the 
number of liters of carbon dioxide exhaled in twenty-four hours by such 
a person. 

23. 100 kg. of coal containing 85 per cent of carbon is burned. 

(a) What weight of carbon dioxide is produced? (b) What volume will 
this occupy under standard conditions ? 

24. Suppose you wish to fill a tank holding 50 liters with carbon 
dioxide, (a) What weight of marble will be required for its preparation ? 

(b) What weight of hydrogen chloride will be needed? 

25. What volume of carbon dioxide will be formed in the combus- 
tion of 100 liters of carbon monoxide? (Suggestion. Note that both of 
the compounds involved are gases ; write the molecular equation and 
solve by inspection.) 

26. 1 g. of coal on combustion gave 3 g. of carbon dioxide, (a) What 
is the weight of the carbon in the 3 g. of carbon dioxide ? (b) What is 
the source of this carbon? (c) Calculate the percentage of carbon in 
the sample of coal burned. 



CHAPTER XII 
VALENCE 

Valence of elements. If we note carefully the formulas of 
the compounds given in the previous chapter, it will be seen 
that the atoms of the various elements differ among themselves 
in respect to the number of other atoms with which they com- 
bine. Taking into account for the present only those formulas 
that contain two kinds of atoms, we see that one atom of chlo- 
rine combines with one atom of hydrogen as shown by the 
formula HC1, while one atom of oxygen combines with two 
of hydrogen (H 2 0). Later we shall find that one atom of 
nitrogen combines with three atoms of hydrogen (NH 3 ), while 
one atom of carbon combines with four of hydrogen (CH 4 ). 
Moreover, when sodium acts on water we have seen (p. 97) 
that one atom of sodium displaces one of hydrogen ; on the 
other hand, when zinc acts on sulfuric acid one atom of the 
metal displaces two of hydrogen (p. 97). 

These observations bring into view an important property 
of the atoms called valence. The valence of an atom is that prop- 
erty which determines how many atoms of any other kind it can 
hold in combination or can displace in a reaction. 

Standard of valence. Since the atoms of neither hydrogen 
nor chlorine combine with more than one atom of another kind 
(to form a compound which contains only two kinds of atoms), 
they are called univalent atoms and serve as a standard for 
the valence of other atoms. Oxygen, sulfur, iron, and zinc, 
whose atoms combine with two of hydrogen or of chlorine, are 
called bivalent. There are other elements whose atoms have 

118 



VALENCE 119 

still higher valences, of three (tervalent), four (quadrivalent), 
and so on up to a valence of eight. 

From all that has been said it might be inferred that in the 
compound CuO the atoms both of copper and of oxygen are uni- 
valent, because there is one of each in the molecule of copper 
oxide. But measured by hydrogen, oxygen is bivalent (H 2 0). 
As we shall see below, copper must be bivalent also. We must 
always keep in mind that oxygen is bivalent when we deduce the 
valence of an element from the formula of its oxide. 

Applications of valence. While it is not possible at this 
point to go very far into the subject of valence, the following 
general principles will be of service : 

1. If two elements which have the same valence combine 
to form a compound, they will combine atom for atom, as shown 
in the following formulas, in which the valence of the atom of 
each element is designated by the figure above the symbol : 

11 2 2 33 4 4 

HC1, HgO, A1N, CSi 

2. If two elements which have different valences combine 
to form a compound, then such numbers of atoms of the two 
elements will combine as will add up an equal number of 
valences, thus: 

12 31 41 42 32 

H 2 0, NH 3 , CH 4 , C0 2 , A1 2 3 

Hydrogen is univalent, while oxygen is bivalent, as is expressed 
by the figures over the formula of water. There is 1 atom of oxygen 
(2 valences) ; hence there must be a sufficient number of hydro- 
gen atoms to add up 2 valences ; namely, 2 atoms. 

Again, in the formula C0 2 the carbon is quadrivalent. Hence, 
for each atom of carbon (4 valences) there must be 2 atoms of 
oxygen (4 valences). 

3. In any reaction in which one element takes the place of 
another in a compound, or one in which one element of a com- 
pound exchanges places with an element in another compound, 



120 CHEMISTRY AND ITS USES 

the exchange must be between such numbers of atoms of the 
two elements involved as will add up an equal number of 
valences ; for example, one atom of a bivalent element will 
displace two atoms of a Univalent element, while two atoms 
of a tervalent element (6 valences) will displace three atoms of 
a bivalent element (6 valences). 



2 Na + H S0 4 >■ NajSO, + H. 

2 4 2 4 5 

Zn + 2 HC1 y ZnCl + H 



NaOH+HCl ^NaCl + HOH, or HO 



Ba0 2 + H 2 S0 4 — *- BaS0 4 + H 2 2 

With very few exceptions the elements do not have a single 
fixed valence. Most of them have two valences and some have 
more. Thus mercury forms the compounds HgCl and HgCl 3 , 
in which the mercury is univalent in the one and bivalent in 
the other. Similarly, compounds represented by the formulas 
SnCl 2 and SnCl 4 are formed by tin. Moreover, an element 
often has one valence for one kind of an element and a differ- 
ent valence for another. Thus sulfur is bivalent in the com- 
pound H 2 S, but quadrivalent in the compound S0 2 . 

Precautions in the use of valence. While the general idea of 
valence is often of great assistance, yet it must be added that 
the student will meet with many cases which (at first sight at 
least) are misleading. Thus, in the formula for hydrogen per- 
oxide, H 2 2 , one would be inclined to say that hydrogen and 
oxygen have the same valence ; while the formula for iron 
oxide, Fe 3 4 , would indicate that the iron has a valence of 
2|, which is an absurdity. It is not possible in an elementary 
book to go into detail in all these cases, but some of them will 
be explained in later chapters. 

Table of valences. The usual valences of the elements will 
become familiar as we study formulas of their compounds. In 



VALENCE 121 

the following table, however, are given the usual valences of 
a few of the most important elements. 

VALENCE OF ELEMENTS 

Univalent H, Na, K, Ag, CI, Br, I 

Bivalent Ca, Ba, Mg, Zn, Hg, Fe, S, O 

Tervalent Al, Bi, As, Sb, Fe, N, P 

Quadrivalent Sn, C, Si, S 

Quinquivalent N, P, As, Sb 

EXERCISES 

Note. To find the symbols for the various elements consult the table of the 
elements. 

1. Assuming that hydrogen is always tervalent and oxygen always 
bivalent, state the valence of each of the elements in the following 
compounds : 

NH 3 , CO, C0 2 , HBr, PH 3 , S0 2 , Bi 2 3 

2. In the common oxide of arsenic, the arsenic is tervalent. What is 
the formula of the oxide ? 

3. Antimony (tervalent) combines directly with sulfur (bivalent). 
What is the formula of the resulting compound ? 

4. Iron (bivalent) reacts with both hydrochloric and sulfuric acids, 
displacing the hydrogen. Write the equation for each of the reactions. 

5. Aluminium (tervalent) reacts with hydrochloric acid, displacing 
the hydrogen. Write the equation for the reaction. 

6. Complete and balance the following equation, in which the hydro- 
gen and mercury (bivalent) exchange places : 

H 2 S + HgCl 2 y 

7. Phosphorus forms two different compounds with chlorine, in 
which the phosphorus is respectively tervalent and quinquivalent. 
What is the formula of each of the compounds? 

8. Phosphorus forms two different compounds with oxygen, in 
which the phosphorus is respectively tervalent and quinquivalent. 
What is the formula of each of the compounds? 



CHAPTER XIII 
THE AIR 

Introduction. The words atmosphere and air are often used in- 
terchangeably, although, strictly speaking, the former is applied 
to the entire gaseous envelope surrounding the earth, while the 
latter is applied to a limited portion, such as the air of a room. 
Like water, air was formerly regarded as an element. Near the 
close of the eighteenth century, however, through the experi- 
ments of Scheele, Priestley, Cavendish, and Lavoisier, it was 
shown to be a mixture of at least two gases — those which we 
now call oxygen and nitrogen. By absorbing the oxygen from 
an inclosed volume of air and measuring the contraction in 
volume due to the removal of oxygen, Cavendish was able to 
determine with considerable accuracy the relative volumes of 
oxygen and nitrogen present. 

Inasmuch as air is composed principally of a mixture of oxy- 
gen and nitrogen, elements which have already been described, 
its properties may be inferred largely from those of the two 
gases. One liter weighs 1.2928 g. 

Composition of the air. The normal constituents of air, 
together with the approximate volumes of each in samples 
collected in the open fields, are as follows : 



Oxygen 
Nitrogen . 
Water vapor . 
Carbon dioxide 
Argon . 
Helium, neon 
Krypton, xenon 



21 volumes in 100 volumes of dry air 

78 volumes in 100 volumes of dry air 

variable within wide limits 

3 to 4 volumes in 10,000 volumes of dry air 

0.940 volumes in 100 volumes of dry air 

traces 

traces 

122 



THE AIR 123 

In addition there are usually present small quantities of hydro- 
gen peroxide, ammonium nitrate, microorganisms, dust particles, 
and traces of hydrogen. Although not definitely proved, it is 
probable that small amounts of ozone are also present. The air 
in large cities and manufacturing districts is also likely to con- 
tain certain gases evolved in manufacturing processes. Among 
these are hydrogen sulfide (H 2 S) and sulfur dioxide (S0 2 ). 

Water vapor in the air. The quantity of water vapor which 
may be present in the air varies with the temperature. This 
is shown in the following table, which gives the weight in 
grams of the water vapor that 1 cu. m. of air can absorb at the 
temperature indicated : 

Temperature, 0° 10° 20° 30° 

Weight of water, 4.8 g. 9.9 g. 17.1 g. 30 g. 

Constituents of the air that are essential to life. The con- 
stituents that are known to be essential to life are oxygen, 
nitrogen, water vapor, and carbon dioxide. The oxygen in the 
atmosphere directly supports life through the process of respi- 
ration. The nitrogen serves to dilute the oxygen and thus to 
diminish the intensity of its action. It is likewise assimilated 
by certain plants (p. 79). The water vapor prevents excessive 
evaporation of the water present in organisms, while the carbon 
dioxide is an essential plant food. 

Quantitative analysis of air. A number of different methods 
have been devised for the determination of the percentages 
of the constituents of the atmosphere. Among these are the 
following : 

1. Determination of oxygen. The oxygen is withdrawn from 
a measured volume of air inclosed in a tube. This is done 
by means of phosphorus. 

To make the determination a graduated tube is filled with water 
and inverted in a vessel of water. A sample of the air to be ana- 
lyzed is then introduced into the tube until it is nearly filled with 



124 



CHEMISTRY AND ITS USES 



the gas, and the volume is carefully noted. A small piece of phos- 
phorus is attached to a wire and brought within the tube as shown 
in Fig. 70. After a few hours the oxygen in the inclosed air will 
have combined with the phosphorus, the water rising to take its 
place. The phosphorus is removed, and the volume is again noted. 
The contraction in the volume of the air is 
equal to the volume of oxygen absorbed. 



oo 



Fig. 70. The deter- 
mination of the per- 
centage of oxygen in 
air by means of 
phosphorus 



2. Determination of nitrogen. If the gas 
left after the removal of oxygen from a 
portion of air is passed over heated magne- 
sium, the nitrogen is withdrawn, leaving 
argon and the other rare elements. It may 
thus be shown that of the 79 volumes of 
gas left after the removal of the oxygen 
from 100 volumes of air, approximately 
78 are nitrogen and 0.94 argon. The other 
elements are present in such small quan- 
tities that they may be neglected. 

3. Determination of water vapor and carbon 
dioxide. These constituents are determined 
by passing a known volume of air through 

two tubes, the first containing calcium chloride and the second 
sodium hydroxide (see arrangement of tubes, Fig. 64). The cal- 
cium chloride removes the moisture, while the sodium hydrox- 
ide removes the carbon dioxide. The increase in the weights 
of these two substances will give the weights of moisture and 
carbon dioxide respectively in the original volume of air. 

Processes tending to change the composition of the air. These 
processes naturally fall into two classes : those which increase 
the carbon dioxide and those which diminish it. 

1. Processes tending to increase the quantity of carbon dioxide. 
Not only do large quantities of carbon dioxide escape into the 
atmosphere from volcanoes and crevices in the earth's crust, 
but certain processes are constantly taking place which are 



THE AIR 



125 



attended by evolution of this gas. Chief among these are the 
following : (a) Respiration. In this process some of the oxygen 
in the inhaled air is absorbed by the blood and carried to all 
parts of the body, where it combines with the carbon of the 
worn-out tissues. The products of oxidation are carried back 
to the lungs and exhaled largely in the form of carbon dioxide. 

(b) Combustion. All the ordinary fuels contain large percent- 
ages of carbon. On burning, this is oxidized to carbon dioxide. 

(c) Decay of organic matter. When organic matter decays in 
the air the carbon present is oxi- 
dized to carbon dioxide. 

2. Processes tending to decrease 
the quantity of carbon dioxide. 
There are two general processes 
which tend to diminish the quan- 
tity of carbon dioxide in the 
atmosphere : 

(a) The action of plants. Plants 
have the power, when growing in 
sunlight, of absorbing carbon diox- 
ide from the air, retaining the 
carbon and returning a portion of 
the oxygen to the air. It is from 
this source that plants obtain their 
entire supply of carbon. 

That plants evolve oxygen in the sunlight may be shown as 
follows : Some freshly gathered green leaves are placed under 
water in the jar A (Fig. 71) and covered with the funnel B, the 
stem of which extends into the graduated tube C, likewise filled 
with water. Bubbles of oxygen escape from the surface of the 
leaves and are caught in the measuring tube C. 

(b) The weathering of rocks. Large quantities of carbon di- 
oxide are being constantly withdrawn from the atmosphere 
through its combination with various rock materials. 




Fig. 71. The liberation of oxy- 
gen from plants exposed to 
sunlight 



126 CHEMISTRY AND ITS USES 

Composition of the air constant. Notwithstanding the changes 
constantly taking place which tend to alter the composition of 
the air, the results of a great many analyses of air collected in 
the open fields show that the percentages of oxygen and nitro- 
gen, as well as of carbon dioxide, are very nearly constant. 
Indeed, so constant is the ratio of oxygen to nitrogen that 
the question has arisen whether air is not a definite chemical 
compound. 

Air a mixture. That the oxygen and nitrogen in the air 
are not combined may be shown in a number of ways, among 
which are the following : 

1. When air dissolves in water it has been found that the 
ratio of oxygen to nitrogen in the dissolved air is no longer 
21 : 78, but more nearly 35 : 65. If air were a chemical com- 
pound, the ratio of oxygen to nitrogen would not be changed 
by solution in water. 

2. A chemical compound in the form of a liquid has a 
definite boiling point at a given pressure. Water, for ex- 
ample, boils at 100° under standard pressure. The boiling 
point of liquid air, on the other hand, gradually rises as the 
liquid boils. 

Why the air has a constant composition. If air is a mixture and 
changes are constantly taking place which tend to modify its 
composition, how, then, do we account for the constancy of com- 
position which the analyses reveal ? This is explained by several 
facts : (1) The changes which are caused by the processes of com- 
bustion, respiration, and decay, on the one hand, and the action 
of plants, on the other, tend to equalize each other. (2) The winds 
keep the air in constant motion and so prevent local changes. 
(3) The volume of air is so vast and the changes which occur are 
so small, compared with the total volume, that they cannot be 
readily detected. (4) Finally, it must be noted that only air col- 
lected in the open fields shows this constancy in composition. The 
air in a poorly ventilated room occupied by a number of people 
rapidly changes in composition. 



THE AIR 



127 



Impure air and ventilation. The difference in the per- 
centages of oxygen, carbon dioxide, and moisture present in 
inhaled and exhaled air are shown in the following table: 



Constituent 


Inhaled Air 


Exhaled Air 


Oxvo'eii . ... 


21.00% 

0.01% 

variable 


16.00% 
4.38% 
saturated 


Carbon dioxide 

Water vapor 



The injurious effects resulting from inadequate ventilation 
seem to be due neither to lack of oxygen nor to the excess 
of carbon dioxide ; rather they are due to high temperature 
and to the presence of an abnormal amount of water vapor, 
both of which conditions are apt to prevail in crowded and 
poorly ventilated rooms. 

One of the most important requirements for the well-being of 
the body is that its temperature be kept almost exactly constant 
(98.6°F.). Any marked deviation from this temperature is a seri- 
ous matter. This temperature is kept up by the heat of combus- 
tion within the body, and if there were no opposing process, bodily 
temperature would exceed the safety point at times of great exer- 
tion or when the external temperature is high. Evaporation of 
water absorbs heat and so prevents undue rise in temperature. 
This evaporation takes place partly in the lungs and partly through 
the skin. \Yhen the air surrounding the body contains an ab- 
normal percentage of water vapor this evaporation is greatly 
diminished and physical discomfort occurs. 

Moreover, when the air is perfectly still that portion of the air 
in contact with the body tends to become saturated with moisture, 
and evaporation diminishes ; hence the relief that comes from 
keeping the air in motion, as with an electric fan. 

In general a moisture content of about 70 per cent of that re- 
quired for saturation is most conducive to comfort. The volume 
of fresh air necessary for good ventilation varies greatly with 
conditions, but it may be said to be about 30 cu. ft. per minute 
for each person present. 



128 



CHEMISTRY AND ITS USES 



Liquid air. Air is now liquefied on a commercial scale by 
the apparatus already described (Fig. 32). Liquid air is essen- 
tially a mixture of liquid oxygen (boiling point, —183°) and 




Fig. 72. Pouring liquid air from a thermos bottle 

liquid nitrogen (boiling point, —195.7°) (Fig. 72). When air 
at ordinary temperatures heats this liquid, the gases pass off 
in the order of their boiling points, and in this way both oxygen 
and nitrogen are prepared for commercial purposes. 

EXERCISES 

1. When oxygen and nitrogen are mixed in the proportion in which 
they exist in the atmosphere, heat is neither evolved nor absorbed in 
the process. What important point does this suggest ? 

2. How does the air in manufacturing districts differ in composition 
from that in the open fields ? 









THE AIR 129 

3. When ice is placed in a vessel containing liquid air the latter 
boils violently. Explain. 

4. Does an electric fan lower the temperature (a) of a room? (b) of 
an individual in the room ? 

5. What is the meaning of the word thermos? 

6. What is the cause of the cloud shown in Fig. 72 ? 

7. Why must air be thoroughly dried before liquefying it? 

8. What are some of the uses of liquid air? 

9. Do you think it would be practicable to use liquid air for 
refrigerating purposes ? 

10. When liquid air boils which passes off first, oxygen or nitrogen? 

11. Could animal life continue without the presence of carbon 
dioxide in the atmosphere ? 

12. Would combustion be more intense in liquid than in gaseous air? 

13. In winter the processes (combustion and respiration) which tend 
to increase the volume of carbon dioxide in the air are very active, 
while the opposite process (the action of growing plants) takes place 
only to a small extent. Should you expect a sample of air collected in the 
winter to contain a larger percentage of carbon dioxide than one collected 
in the summer ? 

14. As a general average each person exhales about 460 liters of 
carbon dioxide daily. Calculate the weight of this volume of the gas. 

15. Taking the volumes of the oxygen and nitrogen in 100 volumes 
of air as 21 and 78, respectively, calculate the percentages of these 
elements present in the air by weight. 

16. Assuming that dry wood contains 40 per cent of carbon, all of 
which originally came from carbon dioxide in the air, what weight 
of carbon dioxide would have to be absorbed by a plant to make 500 g. 
of wood? 

17. A tube containing calcium chloride was found to weigh 30.1293 g. 
A volume of air which weighed 15.2134 g. was passed through, after 
which the weight of the tube was found to be 30.3405 g. Find the 
percentage of moisture present in the air. 



CHAPTER XIV 
SOLUTIONS 

Definitions. Everyone is familiar with the fact that when 
common salt is stirred in water, the salt disappears in the 
liquid. The chemist expresses this fact by saying that the 
salt dissolves in the water, and he calls the resulting product 
a solution. Just as the salt dissolves in water, so many other 
solids dissolve in water or other liquids forming solutions. 
In all such solutions the liquid is known as the solvent, and 
the dissolved solid is known as the solute. Thus in a solution 
of salt in water, the water is the solvent and the salt the solute. 

It often happens that the solid does not really dissolve, but 
remains suspended in the liquid, rendering it cloudy. This is 
true of clay shaken with water. Upon standing quietly for 
some time, however, the matter in suspension gradually settles, 
while matter in true solution does not. 

If we set aside in an uncovered dish any ordinary solution, 
such as a solution of common salt in water, the solvent gradually 
evaporates and the solid is recovered in its original state. The 
process may be hastened, of course, by heating the solution. In 
some cases solution is preceded by chemical action, and in such 
cases the original solid is not recovered upon evaporation. Thus, 
we say that sulfuric acid dissolves zinc. As a matter of fact, how- 
ever, the acid changes the zinc into zinc sulfate, and it is not the 
zinc but the zinc sulfate which dissolves and which is recovered 
upon evaporation. 

Saturated solutions. If we add a little salt to (say) 500 cc. 
of water in a beaker and stir it, it will readily dissolve. If 
we continue to add more of the salt, a little at a time, a 

130 



SOLUTIONS 



131 



point will be reached finally at which the salt apparently no 
longer dissolves but settles to the bottom of the beaker, and 
we say that the solution is saturated. Experiments show that 
the salt really continues to dissolve, but that the rate at which 
the solid dissolves is just balanced by the rate at which it sepa- 
rates again from the solution. This condition may be described 
by saying that the solid and solu- 
tion are in equilibrium with each 
other. A solution is said to be sat- 
urated at a given temperature when 
it remains unchanged in concentration 
in contact with some of the solid. The 
weight of the solid which will com- 
pletely saturate a definite volume of 
the liquid at a given temperature is 
called the solubility of the substance 
at that temperature. It is usually 
expressed by stating the number of 
grams of solid that will dissolve in 
100 cc. or in 1 liter of the liquid. 




Fig. 73. The rapid growth of 
a crystal suspended in a super- 
saturated solution 



Supersaturated solutions. Most solids 
are more soluble in hot than in cold 
liquids, and a liquid saturated at a 
high temperature usually deposits the 

excess of solute in the form of crystals as the temperature falls, 
maintaining saturation at all temperatures. Sometimes the crys- 
tals fail to form as the solution cools, especially if it is not dis- 
turbed in any way. The solution then contains more of the solute 
than is normally present when the solution is in equilibrium with 
the solid. Such a solution is said to be supersaturated. When a 
crystal of the solid is added to the supersaturated solution the 
excess of solute at once crystallizes out. This may be shown in a 
striking way by suspending a small crystal by a thread in a super- 
saturated solution (Fig. 73). The crystal grows rapidly in size, and 
fragments, breaking off, start crystallization at other points. 



132 CHEMISTRY AND ITS USES 

Classes of solutions. All gases mix freely with each other 
in all proportions, and such mixtures may be regarded as 
the solution of one gas in another. Gases, liquids, and solids 
dissolve in liquids, and one solid frequently dissolves in 
another. The most familiar of these classes are solutions of 
gases or solids in liquids. 

Conditions affecting solubility. A number of different con- 
ditions influence the solubility of a substance in a liquid. 

1. Nature of the solute. Each substance has its peculiar 
solubility just as it has its own odor, taste, or crystalline 
form. All substances may be regarded as being to some 
extent soluble in every liquid, but in many cases the solu- 
bility is so small that it cannot be measured. In other cases 
it is very great. Some solids dissolve in less than their own 
weight of water, and some gases, such as ammonia, dissolve 
to the extent of 1000 volumes in 1 volume of water. 

2. Nature of the solvent. The nature of the solvent is no 
less important. Water and alcohol each has its own peculiar 
solvent power. Water is probably the most general solvent 
for all classes of materials, and alcohol is perhaps next to it. 
Ether, chloroform, and benzene are good solvents for organic 
substances such as fats, waxes, and oils, — a property that is 
utilized in removing grease spots from fabrics. 

3. Temperature. The weight of a solid which a given liquid 
can dissolve varies with the temperature. Usually it increases 
rapidly as the temperature rises, so that the boiling liquid 
dissolves several times the weight which the cold liquid will 
dissolve. In some instances, as in the case of the solubility 
of common salt in water, the temperature has little influence, 
and a few solids are more soluble in cold water than in hot. 

In the case of gases, on the other hand, the lower the tem- 
perature of the liquid, the larger the volume of gas which it can 
dissolve. While most gases can be expelled from a liquid by 
boiling the solution, some cannot. 



SOLUTIONS 133 

Tables of solubilities. For convenience of reference the facts 
known about the solubilities of various substances have been col- 
lected into tables of solubilities, and these are constantly used by 
the chemist. Tables giving the solubilities of a few of the most 
familiar substances will be found in the Appendix. 

4. Pressure. Change of pressure has little effect upon the 
solubility of a solid, but greatly influences that of a gas. The 
iveight of a gas which dissolves in a given case is proportional to 
the pressure exerted upon the gas (Henry's law). If the pres- 
sure is doubled, the weight of the gas going into solution is 
doubled ; if the pressure is diminished one half, then but half 
as much gas will dissolve. Under high pressure large quan- 
tities of a gas can be dissolved in a liquid, and when the pres- 
sure is removed the escape of the gas causes the liquid to foam, 
or effervesce, as in the familiar example of soda water. 

Characteristic properties of solutions. A few general state- 
ments may be made in reference to some characteristic 
properties of solutions. 

1. Distribution of the solid in the liquid. A solid, when dis- 
solved, tends to distribute itself uniformly through the liquid, 
so that every part of the solution has the same concentration. 
The process goes on very slowly unless hastened by stirring 
or shaking the solution. 

2. Boiling point of solutions. The boiling point of a liquid 
is raised by the presence of a substance dissolved in it. In 
general the extent to which the boiling point of a solvent 
is raised by a given substance is proportional not to the 
number of grams of the substance dissolved in a definite 
weight of the solvent, but to the number of gram-moleadar 
weights dissolved. Thus the gram-molecular weight of alcohol 
(C 2 H 6 0) is 46, while that of sugar (C 12 H 22 O n ) is 342 ; and 
46 g. of alcohol dissolved in a definite weight of water will 
cause the same rise in the boiling point of the water as will 
342 g. of sugar. 



134 CHEMISTRY AND ITS USES 

3. Freezing point of solutions. The freezing point of a liquid 
is lowered by the presence of a substance dissolved in it. The 
lowering of the freezing point obeys a law similar to the one 
which holds for the raising of the boiling point. 

Electrolysis of solutions. Pure water does not appreciably 
conduct the electric current. If, however, certain compounds 
such as common salt or hydrochloric acid are dissolved in the 
water, the resulting solutions are found to be good conductors. 
Such compounds are called electrolytes. When the current passes 
through a solution of an electrolyte some chemical change always 
takes place. This change is called electrolysis (p. 18). 

The general method used in the electrolysis of a solution 
is illustrated in Fig. 74. Two plates or rods, A and B, made 
of suitable material, are connected with the wires from a bat- 
tery (or dynamo) C and dipped into 
the electrolyte, as shown in the 
figure. These plates or rods are 
called electrodes. The electrode B 
connected with the negative pole 

of the battery is the negative elec- 
Fig. 74. Diagram showing trod Qr cathod while that con . 
the method of electrolysis . 

of a solution nected with the positive pole A is 

the positive electrode, or anode. 
The particular form of apparatus used varies in individual 
cases ; thus in the electrolysis of water the electrodes are 
so arranged (Fig. 9) that the oxygen and hydrogen evolved 
may be collected. 

In this way the electric current is utilized in bringing about 
chemical changes. It is being used more and more in the com- 
mercial preparation of many of the metals and their compounds. 
Thus the metal aluminium, used so largely in the construction 
of automobiles and culinary vessels, is obtained entirely by 
electrolysis. By the same process copper, is purified and various 
objects are plated with gold and silver (electroplating). 




\ 



SOLUTIONS 135 

EXERCISES 

1. Give two examples of common solutions and name the solvent 
and solute in each case. 

2. Distinguish between the following terms : (a) solution, (&) solvent, 
(c) solute, (d) saturated solution, (e) supersaturated solution. 

3. Why does the water from some natural springs effervesce ? 

4. Explain the formation of the small bubbles that form in water 
that is being heated but has not yet reached the boiling point. 

5. Why does not the water of the ocean freeze as easily as fresh 
water ? 

6. Suggest a method for raising the boiling point of water. 

7. Do you see any reason for the common practice of adding salt 
to the water in which potatoes are to be boiled? 

8. Why does a solid dissolve more rapidly when stirred? 

9. Account for the fact that sugar often deposits from sirups, even 
when no evaporation has taken place. 

10. Is it correct to define a saturated solution of any solid in a liquid 
as M one that contains all the solid the liquid will dissolve " ? 

11. (a) Why is alcohol added to the water used in the radiator of 
automobiles in cold weather? (b) Suggest some other substance that 
could be used in place of alcohol. 

12. Wood alcohol (CH 3 OH), grain alcohol (C 2 H 5 OH), and glycerin 
(C 3 H 5 (OH) 3 ) are all soluble in water in all proportions. If you had 
1 kg. of each, which would be the most effective for use to prevent the 
freezing of water in the radiator of an automobile ? 

13. Distinguish between the following terms: (a) electrolyte, (b) elec- 
trolysis, (c) electrode, (d) cathode, (e) anode. 

14. Saturated solutions of each of the solids potassium nitrate, sodium 
chloride (common salt), and calcium hydroxide were prepared by heating 
the solids with water at 100°. The resulting solutions were then cooled 
at 20°. What weight of solid separated in each case ? (See Appendix 
for table of solubilities.) 

15. (a) 10 g. of common salt was dissolved in water and the solution 
evaporated to dryness. What weight of solid was left ? (&) 10 g. of 
zinc was dissolved in sulfuric acid and the solution evaporated to dryness. 
What weight of solid was left? 



CHAPTER XV 

CHLORINE; HYDROGEN CHLORIDE; HYDROCHLORIC ACID; 
ACIDS AND SALTS , 

Introductory. It is important that we should now learn 
something definite concerning the group of compounds known 
as acids. Before taking up the general subject we will first 
study a common and typical member of the group ; namely, 
hydrochloric acid. This acid is a solution of a compound of the 
two elements hydrogen and chlorine. We have already studied 
hydrogen, and it is advisable for us to become acquainted with 
chlorine before proceeding with a study of hydrochloric acid. 

Properties of chlorine. During the World War our papers 
and magazines had much to say concerning chlorine, for it was 
the first substance used by the Germans as a poison gas. Like 
most of the elements so far studied, chlorine is a gas, but, unlike 
them, it has a greenish-yellow color and a peculiar, suffocating 
odor. When inhaled, even in small quantities, it acts injuri- 
ously upon the throat and lungs, while in larger quantities it 
causes death ; hence care must be taken in working with it not 
to breathe it. Chlorine is a heavy gas, being nearly 2.5 times as 
heavy as air. One volume of water under ordinary conditions 
dissolves 3 volumes of the gas. 

Occurrence. Chlorine does not occur free in nature, but its 
compounds are widely distributed. It is found in combination 
with the metals in the form of chlorides, those of sodium, 
magnesium, and potassium being the most abundant. All salt 
water contains these chlorides, particularly the one known 
as sodium chloride (NaCl), or common salt. Very large beds of 
chlorides (Fig. 81) are found in many parts of the world. 

136 






CHLOKIKE 



137 




Historical. Chlorine was discovered by Scheele (p. 53) 
in 1774. It was thought to be a compound, however, until 
1810, when the English chemist Sir Humphry Davy (Fig. 86) 
proved it to be an 
element and gave it Q 

the name which it 
now bears. 

Preparation. Two 
general methods for 
preparing chlorine will 
be described : the lab- 
oratory method and the 
commercial method. 

1. Laboratory method. 
A common method for 
preparing chlorine in 
the laboratory consists 
in warming a mixture 
of hydrochloric acid 
and manganese dioxide (Mn0 2 ). Manganese tetrachloride 
(MnCl 4 ) at first forms, but is unstable and breaks down into 
manganous chloride (MnCl 2 ) and chlorine (Cl 2 ), thus: 

Mn0 2 -f 4 HC1 MVInCl 4 + 2 H 2 

MnCl 4 HVInCl 2 + Cl 2 

The manganese dioxide and the hydrochloric acid are brought 
together in a flask A (Fig. 75), and a gentle heat is applied. The 
chlorine set free escapes through the tube B and is collected in 
the cylinder C by displacement of air, the color showing when 
the cylinder is filled. 

Instead of using hydrochloric acid in the preparation of chlorine, 
it serves just as well to use a mixture of sodium chloride and 
sulfuric acid, since these two react to form hydrochloric acid. In 
this case the complete reaction is expressed by the equation 

2 NaCl + Mn0 2 + 2 H 2 S0 4 KNa a S0 4 + MnS0 4 + 2 H 2 + Cl 2 



Fig. 75. Preparing chlorine by the action of 
hydrochloric acid on manganese dioxide 



138 



CHEMISTRY AND ITS USES 



2. Commercial method. When an electric current is passed 
through an aqueous solution of sodium chloride (Fig. 74), 
chlorine is evolved at the anode and sodium is set free at 
the cathode. The sodium, however, as fast as liberated, 
reacts with the water present, forming hydrogen, which is 
evolved at the cathode, and sodium hydroxide, which remains 
dissolved in the water (p. 32). The products of the reaction, 

therefore, are chlorine, hydrogen, 
and sodium hydroxide. All the 
chlorine prepared for commercial 
purposes in the United States is 
obtained in this way. The method 
has advantages in that sodium 
chloride is cheap and that the 
sodium hydroxide formed in the 
process has many commercial 
uses. The chlorine so obtained 
is either used at the plant for 
bleaching, or is compressed in 
strong iron cylinders and shipped 
in this form, or is passed into 
slaked lime, forming the solid 
known as chloride of lime or 
bleaching powder, which can be easily shipped and from which 
the chlorine can be recovered as needed. 

The apparatus in which the electrolysis of the salt is carried 
out is known as a cell (Fig. 83), and there are several different 
kinds used, each being designated by the name of its inventor. 
The electrodes used are made of carbon. 

Chemical conduct. At ordinary temperatures chlorine is far 
more active chemically than any of the elements we. have so far 
considered ; indeed, it is one of the most active of all elements. 
1. Action on elements. A great many elements combine di- 
rectly with chlorine, especially when hot. A strip of hot copper 




Fig. 76. The burning of pow- 
dered metals in chlorine gas 



CHLORINE 



139 




u 



foil dropped into chlorine burns with incandescence. Antimony 
and arsenic in the form of a fine powder at once burst into 
flame when dropped into jars of chlorine (Fig. 76). 
The products formed in all cases where chlorine 
combines with another element are called chlorides. 
2. Action on hydrogen. Chlo- 
rine has a particularly strong 
affinity for hydrogen, uniting 
with it to form hydrogen chlo- 
ride. A jet of hydrogen burn- 
ing in the air continues to burn 
when introduced into a jar of 
chlorine A (Fig. 77), giving 
a somewhat luminous flame. 
A mixture of the two gases ex- 
plodes violently when a spark is 
passed through it or when it 
is exposed to bright sunlight. 
3. Action on substances containing hydrogen. Not only will 
chlorine combine directly with free hydro- 
gen but it will often abstract the element 
from its compounds. Even water, which 
is a very stable compound, can be de- 
composed by chlorine, the oxygen being 
liberated. This may be shown in the 
following way : 

Action of chlorine on water. A long tube 
of rather large diameter is filled with a con- 
centrated solution of chlorine in water and 
inverted in a vessel of the same solution, as 
shown in Fig. 78, and the apparatus is placed 
in bright sunlight. Very soon bubbles of a 
gas will be observed to rise through the solution and collect in the 
tube, and an examination of this gas will show that it is oxygen. 



Fig. 77. The burning of a jet 

of hydrogen in an atmosphere 

of chlorine 




Fig. 78. Decomposi- 
tion of water by chlo- 
rine in the sunlight 



140 



CHEMISTRY AND ITS USES 



At first there is formed both hydrochloric acid and hypochlorous 
acid (HCIO), but the latter is unstable and breaks down into HC1 

and °a- Cl 2 + H 2 y HC1 + HCIO 

2 HCIO ^2HCl + 2 

4. Action on color substances; bleaching action. If strips of 
brightly colored cloth or some highly colored flowers are 
placed in dry chlorine, no marked change in color is noticed, 

as a rule. If, however, the 
cloth and flowers are first 
moistened, the color rapidly 
disappears, or, in other words, 
the objects are bleached. Evi- 
dently the moisture as well 
as the chlorine is concerned 
in the action. A study of the 
case shows that the chlorine 
sets the oxygen free from the 
water, as shown above, and the 
oxygen so liberated oxidizes 
the color substance (dye), con- 
verting it into a colorless com- 
pound. It is evident from this 
explanation that chlorine will 
bleach only those substances 
which are changed into colorless compounds by oxidation. It 
has no action on such color substances as carbon and hence 
does not affect printer's ink made from carbon. It cannot be 
used for bleaching certain substances like silk and straw, 
since it injures the fabric. 

Fig. 79 illustrates the bleaching action of chlorine. Strips from 
the same piece of cloth are suspended in three jars, of which the 
first contains air, the second dry chlorine, and the third moist 
chlorine. It will be noted that dry chlorine has almost no bleaching 
action, while the moist chlorine has partially removed the color. 




Fig. 79. Bleaching colored cloths by- 
moist chlorine 



CHLORINE 



141 



5. Action as a disinfectant. Chlorine has also marked germi- 
cidal properties, and the free element, as well as compounds 
from which it is easily liberated, are used as disinfectants. 
It is also used to destroy the 
microorganisms in city water 
supplies (p. 65). 

Bleaching. The process known 
as bleaching is an important one 
in connection with many indus- 
tries. Thus the various kinds 
of fabrics woven from vegetable 
fibers, such as flax and cotton, 
are always more or less colored ; 
hence bleaching is necessary if 
a white fabric is desired. This 
was formerly accomplished by 
spreading the cloth on plots of 
grass and exposing it to air and 
sunlight, but the process was 
very slow. The same results are 
now obtained in a short time by 
the use of chlorine. 

Uses. The normal produc- 
tion of chlorine in the United 
States amounts to nearly 500 
tons daily. Most of this is used 
in bleaching fabrics and espe- 
cially in the bleaching of wood 

pulp from which paper is made. It is also used in making 
bleaching powder and such compounds as chloroform (CHC1 3 ) 
and carbon tetrachloride (CC1 4 ). Increasing amounts are now 
compressed in strong steel cylinders called bombs (Fig. 80), 
and in this form shipped to points where it is needed for 
water purification and the preparation of various compounds. 




Fig. 80. Chlorine stored in 

bombs for use in the purification 

of water 



142 CHEMISTRY AND ITS USES 

Chlorine in the World War. Nearly all the poison gases used in 
the World War were either free chlorine or compounds of chlorine. 
Poison gas was first used by the Germans on April 22, 1915, when 
a large quantity of chlorine, previously stored in cylinders hidden 
in the trenches, was allowed to escape. The gas, being heavy, 
clung to the ground and was carried forward by a favorable wind 
(Fig. 82). Later in the war the use of chlorine gave way largely 
to certain of its compounds which are more poisonous than the 
free element. Moreover, in place of trusting to the wind to carry 
the gas, the compounds were filled into shell and then fired into 
the ranks of the enemy. Because of the large demand for poison 
gas, large quantities of chlorine were required. To help meet the 
demand the United States built at Edgewood, Maryland, in 1917- 
1918, a chlorine plant with a daily capacity of one hundred tons, 
— the largest chlorine plant ever constructed (Fig. 83). 

Nascent state. It will be noticed that when oxygen is set 
free from water by chlorine, it is at that instant able to do 
in a short period of time what ordinary oxygen gas cannot do, 
for it bleaches substances which would remain unchanged in 
air or pure oxygen for a long time. It is generally true that 
the activity of an element is greatest at the instant of its libera- 
tion from its compounds. To express this fact elements at the 
instant of liberation are said to be in the nascent state. When 
moist chlorine acts as a bleaching agent it is nascent oxygen 
which does the bleaching. 

Hydrogen chloride (HC1) : properties. Hydrogen chloride is 
a very important compound, since its solution in water con- 
stitutes ordinary hydrochloric acid (sometimes called muriatic 
acid), which has many commercial uses. It is a colorless gas 
and is 1.26 times as heavy as air. When inhaled, it has an 
irritating and suffocating effect. It is very soluble in water, 
1 volume of water, under standard conditions, dissolving 
506 volumes of hydrogen chloride. Hydrogen chloride, when 
exposed to the air, attracts moisture which condenses, forming 
little clouds or fumes. 




Fig. 81. Mining common salt (sodium chloride) for use in the manu- 
facture of chlorine and hydrochloric acid 

Mines of the Retsof Salt Company, near Rochester, New York 




Fig. 82. A view of a chlorine-gas attack made by the French on the 
German trenches in Flanders during the World War 

The chlorine was stored in cylinders placed in trenches, and when the wind 

was favorable the gas was released and was carried forward by the wind 

to the German trenches 




Fig. 83. View of one of the eight-cell rooms in the government 
chlorine plant built at Edgewood, Maryland, during the World War 

This is the largest chlorine plant ever constructed, having a capacity of 
100 tons of chlorine daily. The cells or compartments A, A, in which the 
sodium chloride is decomposed, are arranged side by side. A row of car- 
bon rods D projects into each of the cells and constitutes the cathode, 
while the side of the cell constitutes the anode. A solution of salt enters 
the cells through the pipe B. The electric current enters through wires in C 
and decomposes the salt. The chlorine set free escapes from the cells 
through the pipes E, E into the larger pipe F, F, and is then conducted to 
the liquefying plant where it is liquefied and stored until desired for use. 
The sodium hydroxide formed flows from the cell through the pipes G, G 
into the trough ffand is recovered by evaporating the water 



HYDKOGEN CHLOEIDE 



143 



Preparation. The preparation of hydrogen chloride may be 
described under two general heads : 

1. Laboratory preparation. While hydrogen chloride can be 
prepared by burning hydrogen in chlorine, it is much more 
conveniently obtained by treating common salt (sodium chlo- 
ride) with sulfuric acid. The following equation shows the 
reaction : 2 NaQ1 + H ^ y ^^ + 2 HC1 

The dry salt is placed in the flask A (Fig. 84), sulfuric acid is 
added, and the flask gently warmed. The hydrogen chloride is 
rapidly given off and can be 
collected by displacement of 
air. To prepare a solution of 
the gas the end of the deliv- 
ery tube is fixed just above 
the level of some water in the 
cylinder B. The gas is very 
soluble and is absorbed as fast 
as it escapes from the tube. 

2. Commercial preparation. 

Commercially, hydrogen chlo- 
ride is prepared in connec- 
tion with the manufacture 
of sodium sulfate, the reac- 
tion being the same as that 
just given. It is also pre- 
pared by heating sodium hydrogen sulfate (which is obtained 
in the manufacture of nitric acid) with sodium chloride : 




Fig. 84. The preparation of a solu- 
tion of hydrogen chloride 



NaCl + NaHSO, 



Na 2 S0 4 + HC1 



In either case the hydrogen chloride liberated is passed into 
water, in which it dissolves, the solution forming the hydro- 
chloric acid of commerce. When the materials are pure a color- 
less solution is obtained. The most concentrated solution has 
a density of about 1.2 and contains approximately 40 per cent 



144 



CHEMISTRY AND ITS USES 



by weight of hydrogen chloride. The commercial acid (often 
called muriatic acid) is usually colored yellow by impurities. 

The extreme solubility of hydrogen chloride in water may be 
shown as follows : A perfectly dry flask A (Fig. 85) is filled with 
hydrogen chloride. This flask is connected, by means of a glass 

tube, with a similar flask B, which is 
nearly rilled with water, as shown in 
the figure. The end of the tube open- 
ing into flask A is drawn out to a 
rather fine jet. By blowing into the 
tube C, a few drops of water are 
forced into A. Some of the hydro- 
gen chloride at once dissolves, thus 
diminishing the pressure inside the 
flask. The water then flows contin- 
uously from B into A until nearly all 
the hydrogen chloride is absorbed. 
It is evident that the connection must 
be air-tight. 

Hydrochloric acid : properties and 

Fig. 85. Apparatus to show chemical conduct. While hydrogen 
the extreme solubility of chloride itsel f has but little chem . 
hydrogen chloride in water . . 

ical activity, its aqueous solu- 
tion, namely, hydrochloric acid, has marked properties, the 
most important of which are as follows: 

1. Taste. A dilute solution of the acid has a sour taste 
like that of vinegar. 

2. Action on colored compounds. It acts upon many colored 
compounds and changes their color in some way — very often 
to a red. Thus, it changes the blue color of litmus (a color- 
ing matter obtained from a plant) to red. 

3. Action upon metals and their hydroxides. The hydroxides 
of the metals are compounds of a metal with hydrogen and 
oxygen, such as sodium hydroxide (NaOH). When hydrochlo- 
ric acid is brought in contact with certain metals or with the 




HYDKOCHLOKIC ACID; ACIDS AND SALTS 145 



hydroxide of the metals, the hydrogen of the acid is replaced by 
the metal, thus : RC1 + Na _ > NaC1 + R 

HC1 -f NaOH h NaCl + HOH (H 2 0) 

Acids. Hydrochloric acid is one of a large number of impor- 
tant solutions known collectively as the acids. Thus, in addi- 
tion to hydrochloric acid we have nitric acid, which is a solution 
of the liquid known as hydrogen nitrate (HNO ), and sulfuric 
acid, which is a solution of an 
oily liquid called hydrogen 
sulfate (H 2 S0 4 ). For most 
purposes it is not necessary to 
distinguish in name between 
these compounds and their 
solutions in water, and both 
are* frequently called acids. 

The characteristic proper- 
ties of acids are in the main 
those given above for hydro- 
chloric acid. We may there- 
fore define an acid as a compound of hydrogen whose aqueous 
solution (1) has a sour taste, (JT) changes blue litmus to red, and 
(3) acts upon metals and their hydroxides, forming compounds 
in which the hydrogen of the acid is replaced by the metal. 

Just as the individuals composing a group of people have cer- 
tain characteristics in common and yet may differ widely in certain 
traits, so the acids, although possessing many common properties, 
may differ in other properties. Thus, some are more easily decom- 
posed than others ; some are very poisonous, while others are not. 

Salts. We have seen in the discussion of acids that the 
hydrogen of the acid may be displaced readily by a metal. 
The resulting compound is known as a salt. A salt may there- 
fore be defined as a compound derived from an acid by replacing 
the hydrogen of the acid by a metal. 




Fig. 85 a. The three most common acids 
used in the laboratory 



146 CHEMISTRY AND ITS USES 

Since salts are derived from acids we speak of the salts of the 
acids. Thus, sodium chloride (NaCl) is a salt of hydrochloric 
acid (HC1), while copper sulfate (CuS0 4 ) is a salt of sulfuric 
acid (H 2 S0 4 ). We might expect each acid to form as many salts 
as there are metals. This is not true, however, since some of the 
salts are unstable and have never been prepared. Nevertheless 
the number of salts is very large. 

Radicals and their valence. If we compare the formulas 
H 2 S0 4 , Na 2 S0 4 , and CuS0 4 , it will be noticed that they all 
contain the group of atoms S0 4 . Similarly, the compounds 
represented by the formulas HOH, NaOH, and KOH all 
contain the group of atoms OH, known as the hydroxyl 
group. Groups of atoms which act together as a unit in chemical 
action are called radicals. Radicals do not exist in a free 
state but only in combination with other atoms or radicals. 
Oftentimes a molecule may contain more than one hydroxyl 
or other radical. Thus, calcium hydroxide contains 1 atom 
of calcium and 2 hydroxyl groups in each molecule. In 
such cases it is customary to write the formula Ca(OH) 2 
rather than Ca0 2 H 2 , although either is correct. Similarly, we 
write A1 2 (S0 4 ) 3 rather than A1 2 S 3 12 , and Ca 3 (P0 4 ) o rather 
than Ca 3 P 2 8 . 

The valence of a radical can be determined by noting the 
number of atoms of hydrogen (or some other univalent ele- 
ment) with which it combines. The radical S0 4 is bivalent, 
for it combines with two hydrogen atoms as represented in 
the formula H 2 S0 4 . The radical OH, on the other hand, is 
univalent, for it combines with one atom of hydrogen to form 
water (HOH or H 2 0). 

Formulas of salts. Knowing the valence of a metal and the 
formula of any acid it is easy to deduce the formula of the 
salt which the metal may naturally be expected to form with 
the acid. Thus, suppose we wish to write the formulas of the 
salt which the element calcium forms with the acids HC1, 



HYDROCHLORIC ACID; ACIDS AND SALTS 147 

H 2 S0 4 , and H 3 P0 4 respectively. First write the symbol of 
the element in place of the hydrogen of the acid : 

CaCl CaSO, CaPO. 

4 4 

Now write over the symbols of calcium, chlorine, and each of the 
radicals SO, and PO, the valence of each element or radical : 



2 1 2 2 2 3 

CaCl CaSO, CaPO 



Finally, take such numbers of the two constituents of each 
salt as will add up an equal number of valences. This gives 
the following as the correct formulas : 

CaCl 2 CaS0 4 Ca 3 (P0 4 ) 2 

Naming of acids. The method of naming acids depends 
upon whether the acid consists of two elements or of three. 

1. Binary acids. Acids containing only one element in addi- 
tion to hydrogen are called binary acids. They are given names 
consisting of the prefix hydro-, the name of the second element 
present, and the termination -ic. Examples: hydrochloric acid 
(HC1) and hydrosulfuric acid (H 2 S). 

2. Ternary acids. In addition to the two elements present 
in binary acids, most acids also contain oxygen. These acids 
therefore consist of three elements and are called ternary 
acids. It is the element other than hydrogen and oxygen 
which gives its name to the acid. It usually happens that the 
same three elements can unite hi different proportions to 
make several different acids. The most familiar one of these 
is given a name ending in the suffix -ic, while the one with 
less oxygen is given a similar name, but ends in the suffix -ous. 
Examples : nitric acid (HN0 3 ) and nitrous acid (HN0 2 ). 

In cases where more than two acids are known, use is made of 
prefixes in addition to the suffixes -ic or -ons. Thus the prefix 
per signifies an acid still richer in oxygen ; the prefix hypo sig- 
nifies one with less oxygen. Thus HNO, if it existed, would be 
called hyponitrous acid, while HX0 4 would be called pernitric acid. 



148 



CHEMISTRY AND ITS USES 



Naming of salts. A salt derived from a binary acid is 
given a name consisting of the names of the two elements 
composing it, with the termination -ide. Example: sodium 
chloride (NaCl). All other binary compounds are named in 
the same way. 

A salt of a ternary acid is named in accordance with the 
acid from which it is derived. A ternary acid with the 
termination -ic gives a salt with the name ending in -ate, while 
an acid with the termination -oris gives a salt with the name 
ending in -ite. The following table, giving a number of 
examples, will make the application of these principles clear : 



Acid 


Formula 


Salt 


Formula 


Hydrochloric 

Chlorous 

Chloric 


HCl 

HC10 2 

HCIO3 


Sodium chloride 
Sodium chlorite 
Sodium chlorate 


NaCl 

NaC10 2 

NaC10 3 



Chlorides. Since each acid forms a number of, salts it is 
convenient after describing any acid to note the general 
properties of its salts. Having now studied hydrochloric acid 
we may add that the salts of this acid, namely, the chlorides, 
are with few exceptions solid compounds and are all soluble 
in water except silver chloride and mercurous chloride. Lead 
chloride is soluble in hot water, but not in cold. 



EXERCISES 

1. What is the significance of the terms chlorine and nascent! 
(Consult dictionary.) 

2. Why are laboratory methods of preparation so often different 
from commercial methods? 

3. What factors enter into the cost of any substance such as 
chlorine and hydrochloric acid? 

4. If you were going to build a plant for the commercial prepara- 
tion of chlorine, what factors should you take into consideration in 
selecting a location for the plant? 



HYDROCHLORIC ACID; ACIDS AND SALTS 149 

5. Mention three discoveries made by Scheele. 

6. Why is printer's ink not bleached by chlorine ? 

7. Chlorine has only a comparatively slight color, and this would 
not affect the photographic plate. How, then, account for the cloud 
shown in Fig. 82 ? 

8. Distinguish («) between atoms and radicals ; (b) between acids 
and salts. 

9. Name three foods or drinks that taste sour. How could you 
prove that acids are present in them ? 

10. Name the following compounds and designate to which class 
(acids or salts) each belongs: HBr, HC1, H 2 S, KBr, MgBr 2 , CuS. 

11. How can you determine the valence (a) of atoms ? (b) of radicals ? 

12. The most common of the ternary acids of sulfur is H 2 S0 4 . 
Name the following compounds and designate to which class (acids or 
salts) each belongs: H 2 S0 4 , H 2 S0 3 , H 2 S0 2 , Na 2 S0 4 , ZnS0 4 , Ag 2 S0 3 , 
Na 2 S0 2 . 

13. The most common of the ternary acids of phosphorus is H 3 P0 4 . 
(«) What is its name ? (b) What is the valence of the radical P0 4 ? 
(c) Derive the formula for the salts that you would expect the above 
acid to form with each of the metals sodium, zinc, and aluminium. 

14. Suppose you wished to prepare 5 liters of chlorine in the labora- 
tory. What materials and what weight of each would you require ? 

15. Calculate the percentage of chlorine in sodium chloride. 

16. A plant for the production of chlorine has an output of 18 tons 
of chlorine daily. What weight of sodium chloride would be required 
daily? 

17. (a) What is the weight of 1 liter of concentrated hydrochloric 
acid? (Density = 1.2.) (b) What weight of hydrogen chloride does it 
contain (pp. 143, 144) ? (c) What weight of sodium chloride would be 
required to prepare this weight of hydrogen chloride ? 



CHAPTER XVI 



SODIUM; SODIUM HYDROXIDE; BASES 

Metals and nonmetals. The chemist finds it convenient to 
divide the elements into two general gronps known as the 
metals and the nonmetals. It is the chemical conduct of an 

element that determines to 
which of these two groups it 
belongs. This distinction will 
be described more fully in a 
later chapter. For the present 
it is only necessary for us to re- 
member that all the metals are 
solids (except mercury, which 
is a liquid), that as a rule they 
are good conductors of heat and 
electricity, and (with the excep- 
tion of gold and copper) that 
they have a silvery luster. Most 
of the metals have a high den- 
sity ; a few, such as aluminium 
and magnesium, however, are 
comparatively light ; while three 
(namely, lithium, sodium, and 
potassium) are so light that they will float on water. 

The elements so far studied are all nonmetals. It is ad- 
visable now to study some one metal, and the one known as 
sodium best serves our purpose. 

General discussion. Sodium was first prepared in 1807 by 
Davy (Fig. 86), who isolated it by passing an electric current 

150 




Fig. 86. Sir Humphry Davy 
(1778-1829) 

A distinguished English scientist who 

proved that chlorine is an element 

and first isolated the elements sodium 

and potassium 



SODIUM; SODIUM HYDKOXIDE ; BASES 151 



through melted sodium hydroxide. It is a silver-white metal, 
a little lighter than water, and so soft that it can be molded 
easily by the fingers or pressed into wire. It melts at 97.5°. 
It is very active chemically, combining with most of the non- 
metallic elements, such as oxygen and chlorine. It displaces 
hydrogen from acids (p. 145) as well as from water (p. 31). 
Because of its affinity for oxygen, it tarnishes immediately on 
exposure to air ; hence it is often kept immersed in kerosene, 
since it has no action upon this liquid. 

Occurrence of sodium. Sodium does not occur in nature in 
a free state. The most familiar compound of the element 
found in nature is sodium chloride. 
This is a constituent of all sea waters 
and mineral waters and forms large 
solid deposits in various parts of the 
world (Fig. 81). The element also 
occurs as a constituent of many rocks, 
and its compounds are therefore pres- 
ent in the soil formed by their disinte- 
gration. Other compounds of sodium 
often found in nature are sodium 
nitrate (known commercially as Qhile 
saltpeter), sodium carbonate, and so- 
dium borate, or borax. 

Preparation. Sodium is prepared by 
the same method which led to its dis- 
covery; namely, by passing an elec- 
tric current through melted sodium 
hydroxide. Water must be excluded; otherwise the sodium, 
as fast as it is liberated, will react with the water to form 
sodium hydroxide. 

Technical preparation. The sodium hydroxide is melted in a cy- 
lindrical iron vessel A (Fig. 87). through the bottom of which rises 
the cathode B. The anodes C, several in number, are suspended 




Fig. 87. A Castner cell for 

the electrolytic production 

of metallic sodium 



152 



CHEMISTRY AND ITS USES 



around the cathode from above. A cylindrical vessel E floats 
in the fused alkali, directly over the cathode, and under this 
cap the sodium and hydrogen liberated at the cathode collect. 
The hydrogen escapes by lifting the cover, and the sodium, pro- 
tected from the air by the hydrogen, is from time to time skimmed 
or drained off. Oxygen is set free upon the anode and escapes 
into the air through the opening F without coining into contact 
with the sodium or the hydrogen. 



Uses. Sodium is used in preparing certain of its com- 
pounds. It is also used in making the well-known dye indigo. 
Compounds of sodium. Sodium forms many useful com- 
pounds. One of these, sodium hydroxide, is a typical member 
of that important class of compounds known as bases, and 
it is desirable for us to study its proper- 
ties at this time ; the discussion of the 
other compounds of sodium may well be 
deferred to a later chapter. 

Sodium hydroxide (caustic soda) (NaOH). 
Sodium hydroxide is a brittle, white crys- 
talline substance which rapidly absorbs 
w T ater and carbon dioxide from the air. 
For laboratory purposes it is ordinarily 
sold in the form of sticks resembling sticks 
of white candy (Fig. 88). It is very soluble 
in water, and it is the solution that is com- 
monly used. As the name caustic soda 
indicates, it is a very corrosive substance, 
having a disintegrating action on most 
animal and vegetable tissues. It is used 
in a great many chemical industries, its chief uses being in 
the manufacture of soap and paper and in refining petroleum. 
As a household article it is sold under the name of lye. 

Preparation. Sodium hydroxide is prepared commercially 
by two different processes, as follows : 





-I 


|| SODIUM fj 
M HYDROXY : 

m |J 

I / / , BH 



Fig. 88. The form 
of sodium hydroxide 
used in the laboratory 



SODIUM; SODIUM HYDROXIDE ; BASES 153 

1. In the older process sodium carbonate is treated in solu- 
tion with calcium hydroxide (slaked lime). Calcium carbon- 
ate is precipitated according to the equation 

Na 2 C0 3 + Ca(OH) 2 — ^CaC0 3 + 2 NaOH 

The dilute solution of sodium hydroxide, filtered from the 
insoluble calcium carbonate, is evaporated to dryness ; the 
solid is then melted and poured into molds to solidify. 
Most of the sodium hydroxide is made by this method. 

2. The newer method consists in the electrolysis of a solu- 
tion of sodium chloride as explained under chlorine (p. 138). 
Every chlorine plant is, therefore, a plant for producing sodium 
hydroxide. 

Chemical conduct. Sodium hydroxide has certain properties 
in common with the hydroxides of other metals. These com- 
mon properties are as follows : 

1. Action on litmus. Its action on litmus (as well as on 
other colored substances) is just the opposite of that of acids ; 
that is, it turns red litmus to a blue color. 

2. Action on acids. It reacts with acids to form salts 
and water (p. 145), 

NaOH + HG1 — >■ NaCl + HOH 

Bases ; alkalies. Sodium hydroxide is one of a group of 
compounds, each member of which is composed of a metal 
combined with one or more OH radicals. Thus we have sodium 
hydroxide (NaOH), calcium hydroxide (Ca(OH) 2 ), and copper 
hydroxide (Cu(OH) 2 ). These compounds are known collec- 
tively as bases. Like the acids, their basic properties are devel- 
oped only in the presence of water, but as a rule no distinction 
in name is made between the compound and its solution. There- 
fore a base may be defined as a compound of a metal and one or 
more OH groups. In the presence of water (I) it changes red litmus 
to a blue color and (2~) it reacts with acids to form salts and water. 



154 CHEMISTRY AND ITS USES 

Those hydroxides which have the most highly developed basic 
properties are often called alkalies. The common alkalies are so- 
dium hydroxide, potassium hydroxide, and calcium hydroxide. 

Neutralization. We have seen that when an acid and a base 
are brought together a reaction takes place which consists in 
the interchange of the hydrogen of the acid and the metal of 
the base, forming a salt and water, thus : 

NaOH + HC1 y NaCl + HOH 

The interaction of an acid and a base to form a salt and water 
is known as neutralization, and the acid, and the base are said to 
neutralize each other. It is evident from the above equation that 
both the acid and the base disappear as such in the reaction, 
and in their places we have a salt and water. By evaporation 
of the water the salt may usually be obtained in pure form. 
Neutralization, therefore, serves as a general method for the 
preparation of salts. 

Indicators. We have seen that acids turn blue litmus to a red 
color, while bases reverse the change, turning red back to blue. 
We may employ litmus, therefore, to determine whether a given 
solution is acid or basic. It may also be used to determine whether 
or not a solution is neutral (that is, neither acid nor basic), for such 
solutions will have no effect on either blue or red litmus. In addi- 
tion to litmus, certain other colored compounds may be used for 
determining whether solutions are acid, basic, or neutral ; and all 
these are known collectively as indicators. 

Balancing equations of neutralization. In neutralization the hydro- 
gen of the acid combines with the hydroxyl radical (OH) of the 
base to form water ; hence in balancing equations representing 
neutralization we must take the acid and the base in such pro- 
portions that the number of hydrogen atoms of the acid will be 
the same as the number of hydroxyl radicals of the base, thus : 



HC1 + NaOH ►■ NaCl + H 2 

H 2 S0 4 + 2 NaOH ►■ Na 2 S0 4 + 2 H 2 C 

H 2 S0 4 + Ca(OH) 2 ►■ CaS0 4 + 2 H 2 



SODIUM; SODIUM HYDROXIDE; BASES 155 

EXERCISES 

1. What is the word from which the symbol of sodium is derived? 
(Consult dictionary.) 

2. (a) Is sodium an abundant element? (//) What weight of it is 
in your body ? 

3. The ordinary conception of a metal is a substance that is heavy 
and hard. Is this view correct from a chemical standpoint ? 

4. Are there any metals that are liquids at ordinary temperatures ? 

5. In the commercial preparation of what element do we also obtain 
sodium hydroxide? 

6. Distinguish between (a) acids, (b) salts, (c) bases, and (7/) alkalies. 

7. If you knew that a liquid contained either hydrochloric acid or 
sodium hydroxide, how could you easily decide which is present? 

8. Does sugar (C 12 H 22 O u ) neutralize an acid in foods or simply 
mask it? 

9. If your clothing becomes spotted with acids, how would you try 
to remove these spots ? 

10. Complete and balance the following equations : 

(a) HC1 + Fe(OH) 3 >■ (c) HN0 3 + Zn(OH) 2 >■ 

(/>) H 2 S0 4 + Ca(OH) 2 y (d) H 3 P0 4 + Al(OH) 3 >■ 

11. Some soils are acid in character and will not grow certain crops, 
such as clover and alfalfa, (a) How could you tell if a given soil is acid ? 
(//) If acid, suggest a method for correcting the acidity. 

12. What weight of sodium is present in 100 g. of sodium hydroxide ? 

13. For every ton of chlorine made commercially, how many pounds 
of sodium hydroxide are formed? 

14. What weight of hydrogen chloride is necessary to neutralize 
100 g. of sodium hydroxide ? 

15. What weight of sulfuric acid (hydrogen sulfate) will be neutral- 
ized by 40 g. of sodium hydroxide ? 



CHAPTER XVII 



IONIZATION 



Introductory. Pure water does not appreciably conduct the 
electric current, or, as we say, it is a nonconductor. If certain 
compounds, such as sugar, are dissolved in the water, the solu- 
tions are also nonconductors. If, however, certain other com- 
pounds, such as sodium chloride, hydrogen chloride, or sodium 
hydroxide, are dissolved in the water, the resulting solution is 
found to be a conductor. A compound whose solution conducts 
the electric current is called an electrolyte. It has been found 

by experiment that all acids, 
bases, and salts are electrolytes. 
We have already seen that 
the electric current in passing 
through such a solution causes 
certain chemical changes (elec- 
trolysis, p. 134). 

Fig. 89 illustrates a very con- 
venient apparatus for determin- 
ing whether a solution is a good 
conductor. The solution is placed 
in the bottle A and the electrodes are dipped into it. Connection 
with the lighting circuit is made by the cord and plug B. If the 
solution is a good conductor, the current will flow through the 
lamp C, which will then glow. 

Theory of ionization. In order to explain the facts just 
stated, as well as many others, the Swedish chemist Arrhenius 
(Fig. 90), in 1887, proposed a theory known as the theory of 
ionization. Its main points are as follows: 

156 




Fig. 89. A convenient form of ap- 
paratus for determining whether 
or not a substance is a conductor 
of electricity 



IONIZATION 



157 



It is supposed that the molecules of all electrolytes when 
dissolved in water tend to fall apart into two kinds of atoms, 
or groups of atoms, known as ions, and that it is these ions 
which carry the electric current through the solution. More- 
over, each of the two kinds of ions formed from any molecule 
are highly charged with elec- 
tricity ; these charges, however, 
are equal and opposite in char- 
acter, so that the solution, as 
a whole, is electrically neutral. 
The chemist expresses these 
facts as follows, using NaCl 
and H 2 S0 4 as examples : 

NaCl >- Na + + Cl~ 

H SO d — y 2 H+ + SO " " 

2 4 '4 

It is not to be supposed 
that in every solution all the 
molecules are ionized. The per- 
centage of molecules ionized 
depends (1) upon the compound, 

(2) upon the solvent (water is 
the best ionizing solvent), and 

(3) upon the amount of water 
present, the greater the amount of water present the greater 
the percentage of molecules ionized. In all such solutions the 
molecules are in equilibrium with the ions, some molecules 
falling apart into ions, while other ions are recombining to 
form molecules as expressed in the equation 




Fig. 90. Svante August Arrhenius 
(1859- ) 

A Swedish chemist who suggested 
the theory of ionization 



NaCl 



Na + + CI 



The double arrows are used to indicate the fact that the 
action is going on in both directions, the result being a state 
of equilibrium. 



158 CHEMISTRY AND ITS USES 

Since solutions of salt contain the ions Na + and Cl~, the ques- 
tion might arise as to why such a solution shows no evidence of 
sodium and chlorine. The reason is that sodium and chlorine are 
greatly modified in properties by the high charge of electricity 
which they carry as ions. Thus the sodium ion does not decom- 
pose water ; if, however, we could discharge the ion, it would then 
become an atom of sodium and would at once decompose the water. 
We shall see later how this can be done. 

To be of value the above theory must give a satisfactory 
explanation of the properties of solutions. We shall now see 
if the theory meets this test. 

The theory of ionization and electrolysis. We have learned 
(p. 138) that when an electric current is passed through a solu- 
tion of sodium chloride there are formed hydrogen, chlorine, 

and sodium hydroxide. The 

D theory of ionization accounts 
for the formation of these 
products from the salt in the 
following way : The solution 




■™i|.|i|i|.|i|if; 



c of salt contains the ions Na + 

Fig. 91. A diagram illustrating and Cl~ (Fig. 91). Now Slip- 

the theory of the electrolysis of , ° . ,, 

sodium chloride (NaCl) P ose that we P lace in the 

solution the electrodes A and 

B, connected with the battery C, so that A forms the anode 

and B the cathode. The anode A being positively charged 

attracts the negatively charged chlorine ions (unlike electric 

charges attract each other). The chlorine ions, on reaching 

the anode, give up their charge and then combine to form 

molecules and are evolved as free chlorine. The sodium ions, 

on the other hand, are attracted to the negatively charged 

cathode B. On coming in contact with the cathode they give 

up their charge and, being now ordinary sodium atoms, at 

once decompose the water present, forming hydrogen and 

sodium hydroxide (p. 31). 



IONIZATION 159 

It is to be carefully noted that the current does not bring 
about the decomposition of the electrolyte into ions, but can pass 
through the solution only when ions are already present. 

Acids, bases, neutralization, and salts from the standpoint 
of the theory of ionization. Acids, bases, and salts are all 
electrolytes, hence they are ionized in solutions. 

1. Acids. All acids ionize in solution, forming positively 
charged hydrogen ions, while the remainder of the molecule 
forms the negatively charged ion. The hydrogen ion is com- 
mon to all acids, while the other ion varies with the acid. 

HC1 ^H + +C1" 

Hence the common properties of acids, such as their taste and 
their action on litmus and metals, must be due to the hydro- 
gen ions. From the standpoint of the theory of ionization an 
acid may be defined as a compound which produces hydrogen ions 
when dissolved in water. 

2. Bases. All bases ionize, forming a metal ion and a 
hydroxyl ion, thus : 

NaOH >■ Na + + OH~ 

The metal ion varies according to the base, but the hydroxyl 
ion is common to all ; hence the common properties of bases 
must be due to the hydroxyl ions. A base may therefore 
be defined as a compound which produces hydroxyl ions when 
dissolved in ivater. 

3. Neutralization and salts. Experiments show that hydrogen 
ions and hydroxyl ions, when brought together, at once unite 
to form water, a compound which is at most but slightly ionized. 
If, then, an acid and a base are brought together, the hydrogen 
ions of the acid and the hydroxyl ions of the base unite, thus : 

H + , CI" + Na + , OH- — y Na + , CI" + H 2 

If the water is evaporated, the Na + of the base and the Cl _ 
of the acid gradually unite to form the salt NaCl. From 



160 . CHEMISTRY AND ITS USES 

the standpoint of the ionization theory, therefore, the terms 
neutralization and salt may be denned as follows: 

Neutralization consists in the union of the hydrogen ion of an 
acid with the hydroxyl ion of a base to form water. 

A salt is a compound formed by the union of the negatively 
charged ion of an acid with the positively charged ion of a base. 

Strength of acids and bases. Since acid and basic properties 
are due to hydrogen ions and hydroxyl ions, the acid or base 
which will produce the greatest percentage of these ions at a 
given concentration must be regarded as the strongest repre- 
sentative of its class. The acids and bases described in the 
foregoing paragraphs are all quite strong. In 10 per cent 
solutions about half of the molecules are dissociated into ions, 
and this is also approximately the extent to which most salts 
are ionized at this same concentration. 

Methods of expressing reactions between compounds in solu- 
tion. Chemical equations representing reactions between sub- 
stances in solution may represent the details of the reaction, 
or they may simply indicate the final products formed. Thus, 
if we wish to call attention to the details of the reaction be- 
tween sodium hydroxide and hydrochloric acid in solution, 
representing the ions which take part in the reaction, we write 
the equation as follows : 

Na + , OH" + H + , CI" y Na + , CI" + H 2 

If, on the other hand, we wish simply to represent substances 
taking part in the reaction and the final products formed, we 
write the equation thus: 

NaOH + HC1 y NaCl + H 2 

Displacement (or electrochemical) series. Upon bringing a 
piece of zinc into an acid, zinc passes into solution and 
hydrogen is evolved : 

Zn + 2 HC1 >■ ZnCl 2 + H 2 



IONIZATION 



161 



In like manner, when zinc is placed in a solution of a salt of 
copper, such as the sulfate CuS0 4 , zinc passes into solution, 
and a corresponding quantity of copper is precipitated: 

Zn + CuS0 4 y ZnS0 4 + Cu 

On the other hand, copper has no effect upon a solution of 
zinc sulfate. 

It has been found to be possible to arrange hydrogen and 
the metals in a table in such a way that any element in the 
list will displace any one below it from its salts and will in 
turn be displaced from its salts by any one above it. This 
list is called the displacement series or the electrochemical series. 



DISPLACEMENT (ELECTROCHEMICAL) SERIES 



1. Potassium 

2. Sodium 

3. Lithium 

4. Calcium 

5. Magnesium 

6. Aluminium 



7. Manganese 

8. Zinc 

9. Chromium 

10. Iron 

11. Cobalt 

12. Nickel 



13. Tin 

14. Lead 

15. Hydrogen 

16. Copper 

17. Arsenic 

18. Bismuth 



19. Antimony 

20. Mercury 

21. Silver 

22. Platinum 

23. Gold 



This table enables us to foretell many reactions. For ex- 
ample, from the positions of the two metals we should expect 
magnesium to displace tin from its salts: 



Mg + SnCl, 



MgCl+Sn 



We should not, however, expect iron to displace aluminium. 

It is of especial interest to notice the position of hydrogen in 
the series. All the metals above it will evolve hydrogen from 
acids, while those below it will not. In the latter case, if any 
action takes place it is preceded by oxidation. 

It will be recalled that sodium liberates hydrogen from 
water as well as from acids. All the metals above hydrogen 
do this, though in many cases the action is very slow. 



162 CHEMISTKY AND ITS USES 

EXERCISES 

1. State reasons why chemists believe that the molecules of some 
compounds separate into parts when dissolved, in water while others do 
not. 

2. Distinguish clearly between atoms and ions. 

3. In accordance with the theory of ionization a solution of so- 
dium chloride in water contains sodium ions. Why does not the sodium 
decompose the water ? 

4. What is the difference in composition between a concentrated 
solution of sodium chloride and a dilute one ? 

5. How do we account for the fact that a substance develops acid 
properties only when dissolved in water or some similar liquid? 

6. A solution of hydrogen chloride in the liquid known as benzene 
does not conduct the electric current. Should you expect this solution 
to have acid properties ? 

7. From the standpoint of the ionization theory what is the change 
that takes place in neutralization ? 

8. Tin dissolves in hydrochloric acid. Will hydrogen be evolved in 
the process (see displacement series) ? 

9. What metals cannot be used for preparing hydrogen? 

10. What metals displace hydrogen from water ? What ones have we 
actually used for this purpose ? 

11. Should you expect any relation to exist between the position of 
an element in the displacement series and the probability of its occurrence 
in a free state in nature ? 

12. A strip of iron is placed in a solution of copper sulfate (CuS0 4 ). 
Should you expect any reaction to take place ? 

13. Will either copper or silver react with water (see displacement 
series) ? 

14. Suppose you were to electrolyze a solution of a salt of either 
copper or silver, what would become of the metals ? 

15. From the standpoint of the ionization theory write the equation 
for the reaction between nitric acid and sodium hydroxide. 



CHAPTER XVIII 
COMPOUNDS OF NITROGEN 

Occurrence. The large quantity of nitrogen occurring in the 
atmosphere (p. 122) is practically all in the free state. In the 
materials composing the earth's crust, on the other hand, there 
occur in certain localities considerable deposits of nitrogen 
compounds, especially of sodium nitrate (NaNO„). Small 
quantities of nitrogenous compounds also occur in all produc- 
tive soils. From these soils the nitrogen is taken up by plants 
and. built into complex compounds known as protein matter. 
Animals feeding on these plants assimilate the nitrogenous 
matter, which becomes an essential part of the animal tissue. 

While a great many compounds of nitrogen are known, it 
is desirable at this time to study only a few of the simpler ones. 

Ammonia (NH 3 ) 

Several compounds consisting exclusively of nitrogen 
and hydrogen are known, but only one, ammonia, need be 
considered here. A solution of this compound in water (ammo- 
nia water) is a familiar substance, since it is used in our homes 
both as a medicine and as a cleansing agent. 

Properties. Under ordinary conditions ammonia is a gas 
which is little more than half as heavy as air. It is colorless 
and has a strong, suffocating odor often noticed about decay- 
ing organic matter. It is extremely soluble in water, 1 liter of 
water at 0° and 760 mm. pressure dissolving 1298 liters of the 
gas, and at 20° and the same pressure 710 liters. In dissolving 
such large volumes of the gas the water expands considerably, 

163 



164 CHEMISTRY AND ITS USES 

so that the density of the solution is less than that of water, 
the most concentrated commercial solutions having a density 
of 0.88. Ammonia is easily condensed into a colorless liquid, 
and can now be purchased in this form. 

Preparation of ammonia. Ammonia can be prepared in the 
laboratory by a simple method, but the industrial methods 
require elaborate apparatus. 

1. Laboratory method. In the laboratory ammonia is pre- 
pared from ammonium chloride (NH 4 C1), a white solid 
obtained in the manufacture of coal gas. In this compound 
the group NH 4 acts as a univalent radical and is known as 
ammonium. When ammonium chloride is warmed with sodium 
hydroxide the ammonium radical and sodium change places, 
the reaction being expressed in the following equation: 

NH 4 C1 + NaOH —+■ NaCl -f NH 4 OH 

The ammonium hydroxide (NH 4 OH) so formed is unstable 
and breaks down into water and ammonia : 

NH 4 OH — y NH 3 + H 2 

Calcium hydroxide (Ca(OH) 2 ) is frequently used in place of 
the more expensive sodium hydroxide, the equations being 

2 NH 4 C1 + Ca(OH) 2 — > CaCl 2 + 2 NH 4 OH 
2 NH 4 OH — ►■ 2 H 2 + 2 NH 3 

The ammonium chloride and calcium hydroxide are mixed to- 
gether and placed in a flask A, arranged as shown in Fig. 92. The 
mixture is gently warmed, when ammonia is evolved as a gas and, 
being much lighter than air, is collected in B by displacement of 
air, as shown in the diagram. 

2. Preparation from coal. Ordinary soft coal contains small 
percentages of nitrogen and hydrogen, along with the carbon, 
and when the coal is heated in the absence of air (see coal gas) 
a part of this nitrogen and hydrogen is converted into ammonia. 



COMPOUNDS OF NITROGEN 



165 



The gas is purified, and either dissolved in pure, cold water, 
forming aqua ammonia, or absorbed by acids with which it 
unites to form solid compounds. Nearly all the ammonia used 
in the United States comes from this source. A limited amount 
is prepared from cyanamide, as will be explained later. 

3. The Haber process. When a mixture of nitrogen and hydrogen 
subjected to a pressure of 200 atmospheres is heated to about 500° 
in contact with a suitable 
catalyzer, about 8 per cent 
of the nitrogen present 
combines with hydrogen. 
The resulting ammonia is 
dissolved in water as fast 
as formed, and the re- 
maining gases are again 
conducted over the cata- 
lyzer. The process thus 
becomes continuous, more 
nitrogen and hydrogen be- 
ing introduced as needed. 
The nitrogen used is ob- 
tained from the air, and 
the hydrogen is obtained 
from water. During the 
war the Germans prepared 




Fig. 92. The preparation and collection of 
ammonia in the laboratory 



large quantities of ammonia by this process. Moreover, it is 
probable that this method will come into general use as coal 
becomes scarcer. The method was devised by a German chemist 
named Haber ; hence the name Haber process. 

Chemical conduct. Ammonia will not support combustion 
nor will it burn under ordinary conditions. In an atmosphere 
of oxygen it burns with a feeble, yellowish name. When quite 
dry it is not a very active substance, but when moist it com- 
bines with a great many substances, particularly with acids. 

Uses. Ammonia can be condensed to a liquid by the appli- 
cation of a moderate pressure. If the pressure is removed from 



166 



CHEMISTRY AND ITS USES 



the liquid so obtained, this liquid rapidly passes again into the 
gaseous state, and in so doing absorbs a great deal of heat. 
Advantage is taken of this fact in the manufacture of ice. 
This is one of the main uses of ammonia. Large quantities 
of ammonia are also used in the manufacture of ammonium 
compounds such as ammonium chloride (NH 4 C1). 

The manufacture of ice. Fig. 93 illustrates the method of manu- 
facturing ice. The ammonia gas is liquefied in the pipes A, B by 
means of a compression pump. The pipes lead into a large brine 

tank, a cross section of 
which is shown in the 
figure. Into the brine 
(concentrated solution 
of calcium chloride) con- 
tained in this tank are 
dipped the vessels D, E, 
F, filled with pure water. 
The pressure is removed 
from the liquid ammonia 
as it passes through the 
expansion valve C into 
the pipes immersed in 
the brine, and the heat 
absorbed by the rapid 
evaporation of the liquid 
lowers the temperature of the brine below zero. The water in D, 
E, F is thereby frozen into cakes of ice. The gaseous ammonia 
resulting from the evaporation of the liquid escapes at G and is 
again liquefied, and the process is repeated. 

Ammonium hydroxide (NH 4 0H). The solution of ammonia 
in water, commonly called aqua ammonia or ammonia water, 
has strong basic properties. It turns red litmus blue and 
neutralizes acids, forming salts with them. It seems certain, 
therefore, that when ammonia dissolves in water it combines 
chemically with the water according to the equation 




Fig. 93. Diagram illustrating the principle of 
an ammonia ice machine 



NH. + H.O 



NHOH 

4 




Fig. 94. View of the interior of an ice plant 

The floor is made up of small wooden plates, each one forming the cover 
of one of the iron vessels in which the water is frozen. These vessels are 
immersed to their rim in the hrine tank under the floor. The picture 
shows two of these vessels, A, A, heing withdrawn from the hrine. The 
water in these is frozen and the cakes of ice in them are removed by 
immersing the vessels for a short time in warm water and then inverting 
the vessels. The vessels are then filled with distilled water, again immersed 
in the brine, and left until the water is frozen. The picture also shows two 
vessels B, B, from which the ice has just been removed, being filled with 
water. As the water flows in through the hose at the top, the vessels 
gradually sink into the brine. The tops of two other vessels, C, C, are also 
shown. The water in these is frozen and the ice is ready to be withdrawn 




Fig. 95. View of the interior of a cold-storage plant 

The temperature in a room may he kept very low hy placing ahout the 
room strong iron pipes into which liquid ammonia is forced and allowed 
to vaporize. The vapor is then conducted away, compressed again to the 
liquid state, and then returned to the pipes in the cold-storage room, so 
that the process hecomes continuous. Moisture in the air in such rooms 
freezes as it comes in contact with the pipes, so that the pipes are commonly 
surrounded hy a layer of ice. By regulating the rate of evaporation of the 
liquid ammonia, any desired temperature (within limits) can he obtained. 
Since foods do not spoil when the temperature is kept down, it is possible 
in this way to store foods for a long period and to withdraw them as needed 



COMPOUNDS OF NITROGEN 167 

The resulting compound, NH 4 OH, called ammonium hydrox- 
ide, is a base just as sodium hydroxide is a base. The sepa- 
ration of the pure hydroxide from its solutions has not been 
accomplished, for as the solution becomes concentrated the 
compound decomposes again into ammonia and water : 

NH 4 OH — y NH 3 + H 2 

The solution of ammonia in water, therefore, constitutes a state of 
equilibrium between ammonia and water, on the one hand, and am- 
monium hydroxide, on the other. This condition is conveniently 
expressed in the following way : 

NH 3 -f- H 2 :<=> NH 4 OH 

Aqua ammonia is a good solvent for grease and is a familiar 
household article. 

The ammonium radical. The univalent radical NH 4 plays 
the part of a metal in many chemical reactions, and is called 
ammonium. The ending -ium is given to the name to indicate 
the metallic properties of the substance, since the names of the 
metals in general have that ending. The salts formed by the 
action of the base ammonium hydroxide on acids are called 
ammonium salts. Thus, with hydrochloric acid, ammonium 
chloride (NH 4 C1) is formed, in accordance with the equation 

NH 4 OH + HC1 — * NH 4 C1 + H 2 

Acids of Nitrogen 

The most important of the acids of nitrogen are the following : 
nitric acid (HN0 3 ) and nitrous acid (HN0 2 ). 

Hydrogen nitrate (HN0 3 ). Hydrogen nitrate is very difficult 
to prepare because it is very unstable. It is always obtained 
in aqueous solution, and this solution constitutes ordinary 
nitric acid. 

Nitric acid : properties. The concentrated nitric acid of 
commerce contains about 68 per cent of hydrogen nitrate and 



168 



CHEMISTRY AND ITS USES 



32 per cent of water. Such a solution has a density of 1.4. 
When pure it is a colorless liquid, but it is frequently colored 
slightly brown owing to decomposition products. 

Preparation. Nitric acid is ordinarily prepared in the lab- 
oratory by the action of sulfuric acid on sodium nitrate 
(NaN0 3 ), which occurs in large quantities in Chile. When 

cold, concentrated sul- 
furic acid is poured 
upon sodium nitrate no 
chemical action seems 
to take place. If, how- 
ever, the mixture is 
heated in a retort A 
(Fig. 96), nitric acid 
is given off as a vapor 
and may be easily con- 
densed to a liquid by 
passing the vapor into 
a tube B surrounded 
by cold water. An 
examination of the liquid left in the retort shows that it con- 
tains sodium hydrogen sulfate (NaHS0 4 ), only half of the 
hydrogen of sulfuric acid having been replaced by sodium. 
The reaction may be represented by the equation 




Fig. 96. Preparing nitric acid in the lab- 
oratory by the action of sulfuric acid on 
sodium nitrate 



NaN0 3 + H 2 SO J 



NaHS0 4 + HN0 3 






Commercial preparation of nitric acid. In the United States nitric 
acid is prepared commercially by the same method that is used 
in the laboratory. Fig. 97 illustrates a form of apparatus used. 
Sodium nitrate and sulfuric acid are heated in the iron retort A . 
The resulting acid vapors pass in the direction indicated by the 
arrows and are condensed in the glass tubes B, which are covered 
with cloth kept cool by streams of water. These tubes are inclined 
so that the liquid resulting from the condensation of the vapors 
runs back into C and is drawn off into the large vessel D. 



COMPOUNDS OF NITROGEN 



169 



Preparation of nitric acid from the nitrogen of the air. Since 
the supply of sodium nitrate is gradually diminishing, many 
efforts have been made to find a way of making nitric acid 
from the nitrogen of the air. There are two general ways of 
doing this: (1) The first consists in exposing a mixture of 
nitrogen and oxygen to the action of an electrical discharge. 
As a result cer- 
tain oxides of ni- 
trogen form which 
unite with water 
to form nitric acid. 
This method is 
used in Norway, 
where the water- 
falls make cheap 
electrical energy 
possible. (2) The 
better method con- 
sists in first con- 
verting the nitrogen 

of the air into ammonia. This can be done either by the 
Haber process (p. 165) or by a process which will be described 
later under the heading " Cyanamide." If now the ammonia 
mixed with air is passed over a catalytic agent (platinum is 
used) in a heated tube, nitric acid is formed through the 
oxidizing action of the air, thus: 




Fig. 97. 



A commercial still for the production of 
concentrated nitric acid 



NH„ + 2 O, 



HN0 3 + H 2 



Nitric acid necessary for the manufacture of explosives. During 
the World War the Germans were driven off the seas and so 
could not obtain sodium nitrate (all of which comes from Chile) 
for the preparation of nitric acid. As a result they developed and 
used the ammonia process. The method is still largely used in 
Germany and to a limited extent in the United States and bids 
fair to come into general use at no very distant date. 



170 CHEMISTRY AND ITS USES 

Chemical conduct. The important reactions of nitric acid 
are as follows: 

1. Acid properties. Nitric acid has all the characteristics 
of a strong acid, changing the color of indicators and form- 
ing salts with bases. 

2. Oxidizing action. Nitric acid is not so stable as hydro- 
chloric acid. When heated or when exposed for some time 
to the sunlight it undergoes partial decomposition, forming 
water and the two gases nitrogen dioxide and oxygen. Be- 
cause of this property nitric acid is a good oxidizing agent. 
It acts upon the skin, turning it yellow. 

3. Action on metals. When a metal reacts with nitric acid 
a salt of the acid is always formed. Hydrogen, however, is 
never evolved (except with a very dilute solution of acid) 
as in the case of hydrochloric acid, but in its place an oxide 
of nitrogen, generally nitric oxide (NO), is given off. The 
reason for this is as follows: At first thought we might 
expect that with the metals above hydrogen in the displacement 
series hydrogen would be evolved. If we recall, however, that 
hydrogen is a strong reducing agent and that nitric acid is 
a strong oxidizing agent, then it would seem reasonable that 
any hydrogen set free would not be evolved but, as fast as 
liberated, would act upon the nitric acid, forming water as 
indicated by the equation 

3 H + HN0 3 — >■ 2 H 2 + NO 

This is what actually happens. Metals below hydrogen in 
the displacement series ordinarily do not react with acids. 
Nitric acid, however, being a strong oxidizing agent, oxidizes 
the metals. A complex reaction takes place, however, as indi- 
cated in the following equation, which expresses the reaction 
between copper and nitric acid : 

3 Cu + 8 HN0 8 y 3 Cu(N0 3 ) 2 + 2 NO + 4 HO 



COMPOUNDS OF NITROGEN 171 

Aqua regia. Since nitric acid is a good oxidizing agent 
we might expect it to liberate chlorine from hydrogen chloride, 
and this is found to be the case. A mixture of 1 part of nitric 
acid by volume and 3 parts of hydrochloric acid is called aqua 
regia and is one of the strongest solvents known. It owes its 
solvent poivers not to its acid properties but to the nascent chlorine 
ivhich it liberates. Metals such as gold and platinum, which are 
not soluble in any of the common acids, readily dissolve in aqua 
regia, being converted into chlorides by the nascent chlorine. 

Uses. Nitric acid has countless uses in the industries and 
in chemical laboratories. It is most extensively used in the 
manufacture of explosives of various kinds and of celluloid, 
photographic films, and dyes. 

Salts of nitric acid : nitrates. The salts of nitric acid are 
called nitrates. Many of these salts will be described in the 
study of the metals. They are all soluble in water, and, when 
heated to a high temperature, undergo decomposition. In a 
few cases a nitrate, on being heated, evolves oxygen, forming 

2 NaN0 3 ►■ 2 NaN0 2 -f 2 

In other cases the decomposition goes farther, and the metal 
is left as oxide : 

2 Pb(N0 3 ) 2 — y 2 PbO + 4 N0 2 + 2 

The nitrates are most largely used in the manufacture of 
gunpowder, sulfuric acid, nitric acid, and as a fertilizer. 

Nitrous acid (HN0 2 ). It is an easy matter to obtain sodium 
nitrite (NaN0 2 ) by heating sodium nitrate, as explained in the 
previous paragraph. Now when sodium nitrite is treated with an 
acid, such as sulfuric acid, it is decomposed, and nitrous acid is 

NaN0 2 + H 2 S0 4 >■ NaHS0 4 + HN0 2 

The acid is very unstable, however, and decomposes into water 
and oxides of nitrogen. Sodium nitrite is used in the manufac- 
ture of dyes 



172 



CHEMISTRY AND ITS USES 



Oxides of Nitrogen 

The most important of the oxides of nitrogen are the 
following : 

a colorless gas 

a colorless gas 

a reddish-brown gas 

known only at low temperatures 

a white solid 



Nitrous oxide (N 2 0) . . 
Nitric oxide (NO) . . . 
Nitrogen dioxide (N0 2 ) 
Nitrogen trioxide (N 2 3 ) . 
Nitrogen pentoxide (N 2 5 ) 



Ob 



Nitrous oxide (laughing gas) (N 2 0). This gas is most readily 
prepared by heating the solid known as ammonium nitrate : 

NH 4 N0 3 — >2H 2 0+N 2 

It is colorless, is somewhat soluble in water, and in solution 
has a slightly sweetish taste. When inhaled, it -produces a 
kind of hysteria (hence the name laughing gas} 
and, if taken in large amounts, insensibility to 
pain and unconsciousness. It was the first sub- 
stance to be used as an ansesthetic in surgery, 
and it is still used in minor 
operations, such as those of 
dentistry. 

Nitrous oxide is a very 
energetic oxidizing agent. 
Substances such as car- 
bon, sulfur, iron, and phos- 
phorus burn in it almost as 
brilliantly as in oxygen, 
forming oxides and setting free nitrogen. Evidently the oxy- 
gen in nitrous oxide is not held in very firm combination by 
the nitrogen. 

Nitric oxide (NO). Nitric oxide is most conveniently pre- 
pared by the action of nitric acid upon copper : 

3 Cu + 8 HN0 3 — * 3 Cu(N0 3 ) 2 + 2 NO + 4 H 2 




Fig. 98. The preparation of nitric oxide 
by the action of nitric acid on copper 



COMPOUNDS OF NITROGEN 



173 



The metal is placed in the flask A (Fig. 98), and the acid 
added slowly through the funnel tube B. The gas escapes 
through C and is collected over water. Nitric oxide is a color- 
less gas slightly heavier than air. Unlike nitrous oxide, it 
does not part with its oxygen easily, and burning substances 
introduced into this gas are usually extinguished. 

Nitrogen dioxide (N0 2 ). When nitric oxide comes into con- 
tact with oxygen or air, it at once combines with the oxygen, 
even at ordinary temperatures, forming a reddish-brown gas, 
N0 2 , which is called nitrogen 
dioxide : 



2 NO + O, 



2 NO. 



It is a reddish-brown gas of un- 
pleasant odor and is poisonous 
when inhaled. 



OQ 




Fig. 99. The formation of nitro- 
gen dioxide from nitric oxide 
and oxygen 



To show the formation of nitro- 
gen dioxide from nitric oxide and 
oxygen a tube is filled with the 
oxide, inverted in water, and pure 
oxygen is passed into it, as shown 
in Fig. 99. As each bubble of oxy- 
gen enters, it unites with the nitric 
oxide to form the reddish-brown 

dioxide. In a few minutes the color fades (because of the action 
of water upon the dioxide) and the water slowly rises in the tube. 

Nitrogen tetroxide. At lower temperatures nitrogen dioxide 
becomes paler in color and condenses to a pale-yellow liquid. It 
has been shown that this paler gas has the formula N 2 4 and 
is called nitrogen tetroxide. At ordinary temperatures the gas is 
a mixture of the two, and we may express this relation thus : 



Nitrogen dioxide, 2 N0 2 

high temperatures 



nitrogen tetroxide, N 2 4 

low temperatures 



Acid anhydrides. The oxides NO s (nitrogen trioxide) and 
N 2 5 (nitrogen pent oxide) are rarely prepared and need not 



174 CHEMISTRY AND ITS USES 

be separately described. They bear a very interesting relation 
to the acids of nitrogen. When dissolved in water they com- 
bine with the water, forming acids : 

N 2 °3 + H 2 °— ^2HN0 2 
N 2 5 + H P— >2HN0 3 

Many other oxides act in the same way, combining with water 
to form an acid. Such oxides are called acid anhydrides. 

Cyanogen (C 2 N 2 ) ; hydrogen cyanide (HCN) ; hydrocyanic 
acid. At high temperatures carbon unites with nitrogen to 
form the colorless, very poisonous gas known as cyanogen 
(C 2 N 2 ). With hydrogen and nitrogen it forms hydrogen 
cyanide (HCN), a colorless liquid boiling at 26°. Hydrogen 
cyanide vapor has the odor of peach kernels. It is one of the most 
poisonous compounds known and is used to destroy insects, espe- 
cially those on citrous fruit trees, such as lemon and orange 
trees. When dissolved in water hydrogen cyanide forms a very 
weak acid known as hydrocyanic acid or prussic acid. Its salts 
are called cyanides. Sodium cyanide (NaCN) is a white solid, 
and its solution dissolves gold, hence it is used in extracting 
gold from its ores. The acid as well as its salts are all very 
poisonous. 

EXERCISES 

1. Give the formulas and names of all the gaseous compounds so 
far studied. 

2. (a) What is a radical? (&) Give examples taken from the present 
chapter. 

3. Perfectly dry ammonia gas does not affect litmus paper. Explain. 

4. Why is nitric acid said to be a strong acid and ammonium 
hydroxide a weak base? 

5. Sulfuric acid has a strong affinity for water. Could ammonia be 
dried by passing it through sulfuric acid ? 

6. Why is brine used in making artificial ice ? 

7. Distinguish between the terms ammonium and ammonia.. 



COMPOUNDS OF NITROGEN 175 

8. What is the valence of the ammonium radical? 

9. Write the equations for the reaction taking place when ammo- 
nium hydroxide is neutralized by each of the three acids — hydrochloric, 
nitric, and sulfuric. 

10. What are the properties of ammonia that make it suitable for 
the manufacture of artificial ice ? 

11. It is said that nitric acid is formed in the air during thunder- 
storms. Does this seem probable ? 

12. Discuss the energy changes which take place in the manufacture 
of artificial ice. 

13. A stream of cold water is sprayed over the pipes A, B (Fig. 93), 
as shown in the figure. State the reason for this. 

14. Why cannot nitric acid be used in the preparation of hydrogen, 
just as hydrochloric and sulfuric acids are used? 

15. How could you easily distinguish between nitrous oxide and 
nitric oxide ? 

16. Are the poisonous properties of prussic acid due to its acid 
properties ? 

17. Why is the formula of cyanogen written C 2 N 2 rather than CN? 

18. What compound of a metal is formed when it dissolves (a) in 
nitric acid ? (b) in hydrochloric acid ? (c) in aqua regia ? 

19. («) How many liters of ammonia under standard conditions 
dissolve in 1 liter of water (p. 163) ? (7>) What weight of ammonium 
chloride will be required for its preparation ? 

20. (a) W^hat is the density of the concentrated nitric acid of 
commerce (p. 168) ? (b) What is the weight of 1 liter of this acid? 
(c) What weight of hydrogen nitrate does it contain? 

21. Suppose you wish to prepare 100 kilograms of the concentrated 
nitric acid of commerce. What compounds and what weight of each 
should you require for its preparation ? 

22. A chemist has an order from a dentist for 100 liters of nitrous 
oxide. What weight of ammonium nitrate will be necessary for its 
preparation ? 



CHAPTER XIX 

REVERSIBLE REACTIONS; EQUILIBRIUM 

Reversible reactions. We have met with a number of reac- 
tions which are of a special interest because they can go in 
either direction. Thus, in the Haber process (p. 165) nitrogen 
and hydrogen combine to form ammonia ; but ammonia, on 
the other hand, is decomposed into nitrogen and hydrogen 
when heated. The chemist expresses these facts in the fol- 
lowing way: ^ + 3 H, q=>: 2 NH 8 

Similarly, ammonia and water combine to form ammonium 
hydroxide (p. 167), while the latter compound is easily decom- 
posed by heat into the compounds from which it is formed, 

thus : NH 3 + H 2 +=± NH 4 OH 

Such reactions as the above, which may go in either direction, 
are known as reversible reactions. 

Equilibrium. If we remember that the materials taking part 
in a reaction are made up of great numbers of molecules all 
of which are in rapid motion and are constantly changing 
their relations to each other, it is not difficult to see why 
some molecules should be decomposing while others are form- 
ing. In time, however, a condition will be reached in which the 
changes in the one direction will just offset those in the other. 
The average percentage of each material present will then 
remain unchanged. This condition of affairs is called equi- 
librium. Thus, in the Haber process ammonia, hydrogen, 
and nitrogen come to equilibrium when about 8 per cent of 
the nitrogen is combined to form ammonia. 

176 



REVERSIBLE REACTIONS; EQUILIBRIUM 177 

This principle of equilibrium may be illustrated by the move- 
ments of people in a busy town. In the early morning nearly 
everyone on the street is moving in one direction toward busi- 
ness. Presently an increasing number are moving in the reverse 
direction, and by the middle of the morning about as many move 
in one direction as in the other. The population of the town may 
then be said to have reached equilibrium between those who are 
at home and those at work in the town. 

Mass action. Suppose, when equilibrium has been reached, 
we add an additional quantity of one of the acting substances 
— say hydrogen — in the case just mentioned. This will make 
it easier for the nitrogen to act upon the hydrogen, for the 
two kinds of molecules will now meet more frequently. It 
will not at all affect the rate at which ammonia is decompos- 
ing. The net effect will therefore be to bring about a new 
equilibrium in which a larger percentage of ammonia is pres- 
ent. The effect 'produced by an excess of one of the reacting 
materials is called mass action. 

Changing an equilibrium to a completed reaction. If we 
withdraw the ammonia as fast as it is formed, before it has 
time to decompose, the reaction will go on until either the 
hydrogen or the nitrogen is used up. In the Haber process 
this is accomplished by absorbing the ammonia as fast as it 
is formed in water or dilute acids. 

The point of equilibrium can therefore be changed or the equilib- 
rium converted into a completed reaction by changing the acting 
mass of the substances taking part in the reaction. 

Let us take another example. Ordinary limestone is impure 
calcium carbonate (CaC0 3 ). When heated it decomposes into the 
solid, calcium oxide (CaO), known as lime, and the gas, carbon 
dioxide. If heated in a closed vessel equilibrium is soon reached : 

CaCQ 3 -<— >-CaO + CO a 

If heated in an open vessel, however, the carbon dioxide escapes 
as fast as formed and the reaction goes to completion. 



178 CHEMISTRY AND ITS USES 

Equilibrium in solution. In aqueous solution the molecules 
of an electrolyte keep parting into ions ; while the ions, on 
meeting, recombine to form molecules, the result being an 
equilibrium between the two conditions, thus: 

HN0 3 ^=±H + + N0 3 - 

If we mix two electrolytes the equilibrium reached becomes 
much more complicated, for any positive ion may unite with 
any negative one. Thus, when we mix sodium nitrate and 
sulfuric acid in the preparation of nitric acid, we have present 
the ions Na + , N0 3 ~, H + , and S0 4 ~ ~, together with the mole- 
cules NaN0 3 , Na 2 S0 4 , NaHS0 4 , HN0 3 , and H 2 S0 4 . 

Completion of reactions in solution. The chemist makes use 
of reactions to secure various compounds in pure condition, 
and he wishes the yield to be as large as possible. Reactions 
which stop short of completion and end in an equilibrium are 
not suited to manufacturing purposes unless means can be 
found to break up the condition of equilibrium and bring the 
reaction to a definite conclusion. There are three conditions 
under which this may be accomplished. 

1. A volatile gas may be formed. If the reaction is conducted 
under such conditions that one of the products is a gas insol- 
uble in the solvent, the gas will make its escape as fast as it is 
formed, and this action will continue until one or the other 
of the ions taking part in its formation is used up. 

Thus, when we mix sulfuric acid and sodium nitrate 
(p. 168) no visible reaction takes place. But if we heat the 
mixture above the boiling point of nitric acid, then the nitric 
acid formed in the equilibrium between the H + and the N0 3 ~ 
ions is converted into a gas insoluble in sulfuric acid and 
distills away until the N0 3 ~ ions are used up. We then have 
a completed reaction expressed in the equation 

Na + , N0 3 - + H + , HS0 4 " y Na + , HS0 4 " + HN0 3 (gas) 



REVERSIBLE REACTIONS; EQUILIBRIUM 179 



It is in this way that most acids are prepared. Their salts are 
heated with some acid of high boiling point, usually with sulfuric 
acid, which boils at 338°. 

2. An insoluble solid may 
be formed. When hydrochloric 
acid (HC1) and silver nitrate 
(AgNO g ) are brought together 
in solution we have in addition 
to these two compounds the 
ions H + , CI", Ag + , N0 3 " and the 
new combinations HN0 3 and 
AgCl. One of these, namely, 
silver chloride (AgCl), is insol- 
uble in water ; therefore as fast 
as it is formed, it separates from 
the solution as a curdy white 
precipitate (Fig. 100). The re- 
action therefore continues till 
either the Ag + or the Cl~ is used 
up, the completed equation being 

H + , Cl" + Ag + , N0 3 "- 




Fig. 100. Precipitating silver 
chloride by the action of hydro- 
chloric acid on silver nitrate 

H + , N0 3 - + AgCl (solid) 



3. Two different ions may unite to form an un-ionized molecule. 

When we bring together sodium hydroxide and hydrochloric 
acid in solution, we have the ions H + , Cl~, Na + , and OH". 
The H + ions and the OH~ ions unite to form molecules of 
water which do not again part into ions save to a very slight 
extent. This leaves only the ions of NaCl in solution, the 
equation being 



Na + , OH- + H + , Cr 



Na + , CI" + H 2 



Neutralization is practically a completed reaction because water 
is so little ionised. 



180 CHEMISTRY AND ITS USES 



EXERCISES 

1. If the quantity of any substance (say the water in a pond or the 
money in a bank), remains constant, does this necessarily prove that the 
substance is undergoing no change ? Give other examples to illustrate 
this same principle. 

2. Give examples of equilibrium o # ther than those mentioned in this 
chapter. 

3. Suggest conditions for nraking the reaction represented below go 

in either direction : 

NH 3 + H 2 O^ZZ>:NH 4 OH 

4. In the preparation of nitric acid (p. 168) why do we heat the 
mixture of sodium nitrate and sulfuric acid? 

5. (a) In the preparation of hydrochloric acid is it necessary to 
apply heat ? (b) Why ? 

6. Why is heat unnecessary in preparing carbon dioxide from marble 
and hydrochloric acid? 

7. Suppose you were to heat some water in a strong closed vessel to 
2500°; represent the changes taking place in the water at this temperature 
(p. 67). 

8. What reactions should you expect to take place when solutions 
of silver nitrate and sodium chloride are mixed? 

9. Barium sulfate (BaS0 4 ) is a white insoluble compound much 
used in paints. Supj)Ose you had a supply of barium chloride (BaCl 2 ) 
and wished to convert it into barium sulfate. Suggest two methods 
(different in principle) for doing this. 

10. The use of sulfuric acid in preparing other acids, such as nitric 
and hydrochloric, is based on what property of the acid ? 



CHAPTER XX 
SULFUR AND ITS COMPOUNDS 

Introduction. While sulfur is not to be regarded as one of 
the more abundant elements, yet it has been known from the 
earliest times and, indeed, was one of the favorite substances 
of the alchemists, who thought that all metals were made up 
of mercury and sulfur and that by changing the proportions 
one metal could be converted into another. Biblical writers 
were certainly familiar with sulfur, for we read of " the lake 
burning with fire and brimstone." 

Properties. Sulfur occurs in a variety of forms. The ordi- 
nary form is a pale-yellow crystalline solid without taste and 
with but a faint odor. It is insoluble in water, but is freely 
soluble in a few liquids, notably in carbon disulfide. It melts 
when heated and forms a rather thin, straw-colored liquid. 
As the temperature is raised, this liquid turns darker in color 
and becomes thicker, until at about 235° it is almost black 
and is so thick that the vessel containing it can be inverted 
without danger of the liquid's running out. At higher tem- 
peratures it becomes thin once more and boils at 444.6°, form- 
ing a yellowish vapor. When the vapor cools the same changes 
take place in reverse order. 

Occurrence. Sulfur is widely distributed in nature and 
occurs in large quantities in the uncombined form, especially 
in the neighborhood of volcanoes. Sicily has long been famous 
for its sulfur mines, and in more recent years large deposits 
have been found in Louisiana and Texas. It is occasionally 
found in well-formed crystals (Fig. 101). 

181 



182 



CHEMISTRY AND ITS USES 



In combination, sulfur occurs abundantly combined with 
metals ; in smaller amounts it is found in a great variety of 

minerals and is a constituent 
of many vegetable and animal 
substances, especially of the 
yolk of eggs. 

Extraction of sulfur. In 
Louisiana and Texas the 
sulfur occurs in deposits far 
underground and is covered 
with quicksand so that it 
cannot be mined. One of 
these deposits lies at a depth 
of 700 feet, is circular in 
shape, and is about half a 
mile in diameter and 500 feet in thickness. Wells are drilled 
into the deposit, and superheated water (above 160°) is forced 
down through suitable pipes. The hot water not only melts the 
sulfur but, being under pressure, forces the molten sulfur to the 




Fig. 101. Crystals of rhombic sulfur 
as they are found in nature 




Fig. 102. Forcing liquid sulfur from deep wells in Louisiana by means of 
hot water under pressure 



earth's surface through pipes suitably arranged (J, Fig. 102). 
The liquid sulfur then solidifies in very large blocks. A single 
well has produced 500 tons daily, and the product is 99.5 per 



SULFUR AND ITS COMPOUNDS 



183 




Fig. 103. A sulfur still 



cent pure. Practically all the sulfur used in the United 

States comes from the Louisiana and Texas deposits. 

Distillation of sulfur. Sul- 
fur may be distilled by heat- 
ing it in a retort-shaped vessel 
A (Fig. 103), the exit tube of 
which opens into a cooling 
chamber B of brickwork. 

When the sulfur vapor first 
enters the cold chamber it 
condenses as a fine crystalline 

powder called flowers of sulfur (Fig. 104). As the condensing 

chamber becomes warm the sulfur condenses as a liquid and 

is drawn off into cylindrical molds, the product being called 

roll sulfur or brimstone (Fig. 104). 

Varieties of sulfur. Sulfur exists in a number of quite 

different (allotropic) forms. Many other elements occur in a 

number of different forms, but 

those of sulfur are unusually 

numerous and are easy to obtain. 

The best-known are the following : 

1. Ordinary, or rhombic, sulfur. 
When sulfur crystallizes from 
solution in liquids (notably from 
carbon disulfide) it is obtained 
in compact yellow crystals which 
melt at 112.8°. This is called 
rhombic sulfur (Fig. 101), and 
brimstone is composed largely of 
this variety. 

2. Prismatic, or monoclinic, sul- 
fur. When melted sulfur is allowed to cool until a part of 
the liquid has solidified, and the remaining liquid is then 
poured off, it is found that the solid sulfur remaining in the 




Fig. 104. 



Two common forms 
of sulfur 



184 



CHEMISTRY AND ITS USES 



vessel is in the form of fine needle-shaped crystals which 
melt at 119.2°. The needle-shaped form is called monoclinic 
sulfur. At all temperatures below 96° the needle-shaped 
crystals break up more or less rapidly into little crystals of 
the rhombic variety. 

3. Plastic sulfur. When boiling sulfur is poured into cold 
water it assumes a gummy, doughlike form which is quite 
elastic. It is simply undercooled liquid sulfur. It is easily 

made by distilling sulfur from a 
small, short-necked retort (such as 
is represented in Fig. 105) and 
allowing the liquid to run directly 
into water. In a few days it passes 
over into ordinary rhombic sulfur. 
Chemical conduct of sulfur. Sul- 
fur burns in oxygen or in the air 
with a pale-blue flame, forming 
sulfur dioxide (S0 2 ). Most metals 
when heated with sulfur combine 
directly with it, forming metallic 
sulfides. In some cases the action 
is so energetic that the mass becomes incandescent, as is the 
case with iron. This conduct recalls the action of oxygen 
upon metals, and in general the metals which combine readily 
with oxygen are apt to combine with sulfur. 

Uses of sulfur. Large quantities of sulfur are used in the 
manufacture of gunpowder, vulcanized rubber, sulfuric acid, 
and other compounds of the element. It is also used exten- 
sively in the manufacture of insecticides for use in orchards 
and vineyards. 

Lime-sulfur spray. The chief sulfur insecticide is known 
as lime-sulfur spray. It is made by boiling sulfur with slaked 
lime, by which process a deep-red liquor is obtained, consisting 
essentially of a solution of sulfides of calcium (CaS 4 and CaS 5 ). 




Fig 



105. The preparation of 
plastic sulfur 



SULFUR AND ITS COMPOUNDS 



185 



The liquid is a very efficient insecticide, particularly for scale, 
and it is also a fungicide. Large quantities of it are used for 
spraying fruit trees (Fig. 106). 

Hydrogen sulfide (H 2 S) : properties. Sulfur forms with 
hydrogen the important compound known as hydrogen sulfide 
(H 2 S). This is a colorless gas having a weak, disagreeable taste 
and a most offensive odor, suggesting rotten eggs. At ordinary 



W; 








. -^ 


:' V :.*V U~ 


1 


1 


-\ V 7 


?| 


■ 


> & 


^nJiMiWi 




fiS 






$mm'i?$m 




m 


#§ri~ : ""---^ : . ; " jj|. 


- 




T- '■-■- . M 


J &*' ..- 


— . — . -i 



Fig. 106. Spraying an orchard of fruit trees with lime-sulfur spray 

temperatures it is but sparingly soluble in water ; in boiling 
water it is not soluble at all. When inhaled in concentrated 
form it acts as a violent poison, and even when much diluted 
with air produces headache, dizzhiess, and nausea. It is 1.18 
times as heavy as air. 

Occurrence. Hydrogen sulfide is found in the vapors issu- 
ing from volcanoes. It also occurs in solution in many natu- 
ral waters (sulfur waters). It is formed when organic matter 
containing sulfur undergoes decay, just as ammonia is formed 
from nitrogenous matter. 



186 



CHEMISTKY AND ITS USES 




Preparation. Since hydrogen sulfide is a gas which is but 
little soluble in water, it can be prepared by treating a sulfide 
with an acid (p. 179). Iron sulfide (FeS) is usually employed : 

FeS + 2 HC1 — >■ FeCl 2 + H 2 S 

A convenient apparatus is shown in Fig. 107. A few lumps of 
iron sulfide are placed in the bottle A and dilute acid is added a 

little at a time through the funnel tube 
B, the gas escaping through the tube C. 

Chemical conduct. The most im- 
portant chemical properties of hydro- 
gen sulfide are the following: 

1. Acid properties. When dissolved 
in water, hydrogen sulfide acts as a 
weak acid, the solution being some- 
times called hydro sulfuric acid. The 
solution turns blue litmus red and 

Fig. 107. The preparation neutralizes bases, forming salts called 
of hydrogen sulfide by the sulfides. 

2. Action with oxygen. The ele- 
ments composing hydrogen sulfide 

have each a strong tendency to combine with oxygen and are 
not held together very firmly. Consequently the gas burns 
readily in oxygen or air according to the equation 

2 H 2 S + 3 2 — >■ 2 H 2 + 2 S0 2 

When there is not enough oxygen for both the sulfur and 
the hydrogen or when the temperature is kept low, the latter 
element combines with the oxygen, and the sulfur is set free : 

2 H 2 S + 2 — h 2 H 2 + 2 S 

3. Reducing action. Owing to the ease with which hydrogen 
sulfide decomposes, and the strong tendency of both sulfur 
and hydrogen to combine with oxygen, the substance is a 
strong reducing agent. 



action of hydrochloric acid 
on sulfide of iron 



SULFUR AND ITS COMPOUNDS 



187 



4. Action on metals. Hydrogen sulfide acts upon many metals, 
forming sulfides. Silver sulfide (Ag 2 S) is black, and it is owing 
to traces of hydrogen sulfide in the air that silver objects tarnish. 

Sulfur waters. The waters of many natural springs hold 
hydrogen sulfide in solution, as is indicated by their strong odor 
and the way in which they will blacken a silver coin. When the 
water reaches the air the hydrogen sulfide is slowly oxidized, 
with the liberation of sulfur, 
which often deposits about 
the borders of the spring. 

Salts of hydrosulfuric 
acid ; sulfides. The salts of 
hydrosulfuric acid (called 
sulfides) form an important 
class of compounds. Many 
of them are found abun- 
dantly in nature, and some 
of them are important ores. 

Uses of the sulfides in analy- 
sis. Most of the sulfides are 
insoluble in water, and some 
of them are insoluble in acids. Consequently, when hydrogen 
sulfide is bubbled through a solution of a salt, it often happens 
that a sulfide is precipitated. With copper chloride the equation is 

CuCl 2 + H 2 S - — > CuS + 2 HC1 

Because of the fact that some metals are precipitated in this 
way as sulfides, while others are not, hydrogen sulfide is exten- 
sively used in the separation of the metals in the laboratory. 

Laboratory preparation of sulfides. In the laboratory the sulfides 
are prepared by passing hydrogen sulfide through solutions of 
the salts of the metals as shown in Tig. 108. The hydrogen 
sulfide generated in flask A is passed through bottles B and C, 
containing, say, solutions of silver nitrate and arsenic chloride 
respectively. As the gas bubbles through the solutions there is 
formed black silver sulfide in B and yellow arsenic sulfide in C. 




Fig. 108. The preparation of insoluble 
sulfides in the laboratory 



188 CHEMISTRY AND ITS USES 

Oxides of Sulfur 

Sulfur forms two well-known compounds with oxygen: 
sulfur dioxide (S0 2 ), sometimes called sulfurous anhydride ; 
and sulfur trioxide (SO g ), frequently called sulfuric anhydride. 

Sulfur dioxide : properties. Sulfur dioxide is a colorless gas 
which at ordinary temperatures is 2.2 times as heavy as air. 
It has a peculiar, irritating odor. The gas is very soluble in 
water, 1 volume of water dissolving 80 volumes of the gas 
under standard conditions. It is easily condensed to a color- 
less liquid and can be purchased in this condition, stored in 
strong bottles or in metal cylinders. 

Occurrence. Sulfur dioxide often occurs in nature in the 
gases issuing from volcanoes, and in solution in the water 
of many springs. It is likely to be found wherever sulfur 
compounds are undergoing oxidation. 

Preparation. Two general ways may be mentioned for the 
preparation of sulfur dioxide: 

1. By the combustion of sulfur. Sulfur dioxide is readily 
formed when sulfur or certain compounds containing sulfur, 
such as the metal sulfides, are heated in air or oxygen: 

s + o 2 — >so 2 

2 ZnS + 3 2 y 2 ZnO + 2 S0 2 

2. By the reduction of sulfuric acid. When concentrated sul- 
furic acid is heated with certain metals, such as copper, part 
of the acid is changed into copper sulfate and part is reduced 
to sulfurous acid, the latter decomposing into sulfur dioxide 
and water. The details of the reaction will be given later, 
but the complete equation is as follows: 

Cu + 2 H 2 S0 4 — ►» CuS0 4 + S0 2 + 2 H 2 

Chemical conduct. Sulfur dioxide has a marked tendency 
to combine with other substances and is therefore an active sub- 
stance chemically. It combines with oxygen gas, but not very 



SULFUR AND ITS COMPOUNDS 



189 



easily. It can, however, take oxygen away from some other 
substances and is therefore a good reducing agent. Its most 
marked chemical property is its ability to combine with water. 
Sulfurous acid (H 2 S0 3 ). When sulfur dioxide is passed into 
water it combines chemically with it to form sulfurous acid 
(H 2 S0 3 ). It is impossible to prepare this acid in pure form, 
as it breaks down very easily into water and sulfur dioxide. 
The reaction is therefore reversible and is expressed by 
the equation 



HO + SO, 



H£0, 



Solutions of the acid in water are often prepared and have 
a number of interesting properties and commercial uses. 

1. Acid properties. The solution has all the properties of a 
very weak acid. When neutralized by bases, sulfurous acid 
yields salts called sulfites, most of which are insoluble in water. 

2. Reducing properties. Solutions of sul- 
furous acid act as good reducing agents. 
This is due to the fact that sulfurous acid 
has the power of taking up oxygen from 
the air or from substances rich in oxygen, 
and is changed by this reaction into sulfuric 

acid : 2 H 2 S0 3 + 2 — ^2 H 2 S0 4 

3. Bleaching properties. Sulfurous acid 
has bleaching properties, although it is not 
as good a bleaching agent as chlorine 
(p. 140). It is used, however, to bleach 

substances that would be injured by the action of chlorine, 
such as paper, straw goods, and even such foods as canned corn 
and dried fruits. As a rule the bleaching is not permanent. 

The bleaching properties of sulfurous acid may be shown by 
bringing a small dish of burning sulfur under a bell jar (Fig. 109) 
in which has been placed some highly colored flower moistened with 
water. Straw hats may be cleaned and brightened in a similar way. 




Fig. 109. Bleaching 

a red flower with 

sulfurous acid 



190 CHEMISTRY AND ITS USES 

Uses of sulfurous acid. Sulfurous acid is used as a bleach- 
ing agent, as noted above. It also is used in certain foods and 
sirups, serving not only as a bleaching agent but also as a pre- 
servative. Whether or not its use in foods should be permitted 
has been a much-debated question. 

Salts of sulfurous acid ; sulfites. The sulfites are solid 
compounds and, like sulfurous acid, have the power of tak- 
ing up oxygen very readily and are good reducing agents. 
On account of this tendency commercial sulfites are often 
contaminated with sulfates. 

Sulfur trioxide (S0 3 ) : properties. Sulfur trioxide is a color- 
less liquid which solidifies at about 15° and boils at 46°. 
A trace of moisture causes it to solidify into a mass of silky 
white crystals which have the formula S 2 6 . In contact with 
the air it fumes strongly, and when thrown upon water it 
dissolves with a hissing sound and the liberation of a great 
deal of heat. The product of this reaction is sulfuric acid, so 
that sulfur trioxide is the anhydride of that acid (p. 174) : 

SO s + H s O— kH 2 S0 4 

Preparation. When sulfur dioxide and oxygen are heated 
together at a rather high temperature, a small amount of sul- 
fur trioxide ($0 3 ) is formed, but the reaction is slow and 
incomplete. If, however, a suitable catalytic agent is present, 
the reaction is rapid and nearly complete : 

2S0 2 + O a -+2S0 3 

Experimental preparation of sulfur trioxide. The experiment can 
be performed by the use of the apparatus shown in Fig. 110. The 
catalytic agent (fine platinum) is secured by moistening asbestos 
fiber with a solution of chloroplatinic acid and igniting it in a 
flame. The fiber, covered with fine platinum, is placed in a tube 
of hard glass A , which is then heated with a burner to about 400°, 
while sulfur dioxide and air are passed into the tube through the 
drying bottles B and C. Union takes place at once, and the 



SULFUR AND ITS COMPOUNDS 



191 



strongly fuming sulfur trioxide escapes from the jet at the end 
of the tube, and may be condensed by surrounding the receiving 
tube D with a freezing mixture. 

Sulfuric acid (oil of vitriol) (H 2 S0 4 ). Strictly speaking, sul- 
furic acid is a solution of hydrogen sulfate (H 2 S0 4 ) in water. 
It is customary, however, to apply the term sulfuric acid to 
both the hydrogen sulfate and its aqueous solution. The pure 
compound is a colorless, oily liquid nearly twice as heavy as 



SO- 




S> 



~=*^ ^^ ^ 




Fig. 110. The preparation of sulfur trioxide 

water. The ordinary concentrated acid contains about 2 per 
cent of water, has a density of 1.84, and boils at 338°. It is some- 
times called oil of vitriol, since it was formerly made by distilling 
a mixture of substances, one of which was called green vitriol. 
Sulfuric acid is one of the most important of all manufac- 
tured chemicals. Indeed, it has been said that the state of 
civilization of a country can be told by the quantity of sul- 
furic acid used, but such a statement, while probably true, 
could be just as well made in regard to a number of sub- 
stances, such as sodium hydroxide or soap. Enormous quan- 
tities of the acid are used in many of the industries, especially 
in dissolving off the scale from metals, hi the refining of 
petroleum, and hi the manufacture of explosives, dyes, hydro- 
chloric and nitric acids, sodium carbonate, and phosphate 
fertilizers. It is one of the most common laboratory reagents. 



192 CHEMISTRY AND ITS USES 

Manufacture of sulfuric acid. Sulfuric acid can be made at 
low cost and is the cheapest of the commercial acids. Two 
general methods are used in its manufacture : 

1. Contact process. In this process sulfur trioxide is made 
from sulfur dioxide and oxygen, as explained under Fig. 110. 
The two gases are conducted through iron tubes filled with 
some porous material, such as asbestos or sodium sulfate, 
through which is interspersed a suitable catalyzer, such as iron 
oxide or platinum. The sulfur trioxide so formed reacts with 
water to form sulfuric acid: 

S0 3 + H 2 0— ^H 2 S0 4 

The contact process is used only when the concentrated acid 
is desired. 

2. Chamber process. The older method of manufacture, ex- 
clusively employed until recent years and still the most im- 
portant process, is much more complicated. The conversion 
of water, sulfur dioxide, and oxygen into sulfuric acid is 
accomplished by the catalytic action of oxides of nitrogen. 
Since these oxides are gases it is difficult to prevent their 
escape, and very elaborate precautions have to be taken to 
reduce the loss as much as possible. The reactions are brought 
about in very large lead chambers into which oxides of nitro- 
gen, sulfur dioxide, steam, and air (to furnish oxygen) are 
introduced in suitable proportions. 

The dilute acid which collects upon the floor of the lead 
chambers contains from 62 to 70 per cent of hydrogen sulfate. 
It is drawn off and in this form serves for many purposes, 
such as the manufacture of fertilizers. The pure concentrated 
acid can be prepared from the dilute acid by distillation, but 
the process is costly, so that it is sometimes cheaper to prepare 
this form of acid by the contact process. 

Chemical conduct. Sulfuric acid possesses properties which 
make it one of the most important of chemical substances. 



SULFUR AND ITS COMPOUNDS 193 

1. Action as an acid. In dilute solution sulfuric acid acts 
as a very strong acid, turning blue litmus red and forming 
salts with oxides and hydroxides. 

2. Action as an oxidizing agent. Sulfuric acid contains a 
large percentage of oxygen and, like nitric acid, is a very 
good oxidizing agent. When the concentrated acid is heated 
with sulfur or carbon or various other substances, oxidation 
takes place, the sulfuric acid decomposing according to the 
equation H ^ ^ H ^ + Q 

3. Action on metals. Dilute sulfuric acid acts on the metals 
above hydrogen in the displacement series, forming salts of 
the acid (sulfates') and liberating hydrogen (p. 33). The con- 
centrated acid, being a good oxidizing agent, acts in a different 
way upon the metals above hydrogen in the displacement series, 
as well as upon most of the metals below hydrogen. The metal 
is first oxidized by the acid ; the resulting oxide then reacts 
with more of the acid, forming a salt and water. Thus, with 
copper the equations are 

(1) Cu + H 2 S0 4 — >■ CuO + H 2 + S0 2 

(2) CuO + H 2 S0 4 — ^CuS0 4 + H 2 

By canceling the formula of copper oxide in the above equa- 
tions (since the oxide is produced in the first reaction and con- 
sumed in the second) we may combine the two equations into 
a single one ; namely, 

Cu + 2H 2 S0 4 >■ CuS0 4 + 2H 2 + S0 2 

This equation tells us what materials are used up and what the 
final products are ; it does not tell us what materials may have 
been produced and consumed in the process. 

4. Action on salts. We have repeatedly seen that an acid 
of high boiling point heated with the salt of some acid of 
lower boiling point will drive out the low-boiling acid (p. 179). 



194 CHEMISTRY AND ITS USES 

The boiling point of sulfuric acid (338°) is higher than that 
of almost any common acid; hence it is largely used in the 
preparation of other acids. 

5. Action on water. Concentrated sulfuric acid has a very 
great affinity for water and is therefore an effective drying, 
or dehydrating, agent. Gases which have no chemical action 
upon sulfuric acid can be freed from water vapor by bubbling 
them through the concentrated acid. 

Not only can sulfuric acid absorb water but it will often with- 
draw the elements hydrogen and oxygen from a compound con- 
taining them, decomposing the compound and combining with the 
water so formed. For this reason most organic substances, such 
as sugar, wood, cotton and woolen fiber, and even flesh, all of 
which contain much oxygen and hydrogen in addition to carbon, 
are charred by the action of the concentrated acid. 

Salts of sulfuric acid ; sulfates. The sulfates form a very 
important class of salts, and many of them have commercial 
uses. Copperas (iron sulfate), blue vitriol (copper sulfate), and 
Epsom salt (magnesium sulfate) serve as examples. Many 
sulfates are important minerals, prominent among these being 
gypsum (calcium sulfate) and barite (barium sulfate). 

Monobasic and dibasic acids. Acids like hydrochloric and 
nitric acids, which have only one replaceable hydrogen atom 
in the molecule (or, in other words, which yield one hydrogen 
ion in solution), are called monobasic acids. Acids like sulfuric 
acid, which have two replaceable hydrogen atoms in each 
molecule, are called dibasic acids. Similarly, we may have 
tribasic and tetrabasic acids. 

Normal and acid salts. It is possible for such acids as HJS, 
H 2 S0 3 , and H 2 S0 4 to form two kinds of salts. In the one all 
the hydrogen of the acid has been replaced by a metal, as in the 
salts Na„S and Na SO„. These are called normal salts. In the 

2 2 4 

other only one half of the hydrogen has been replaced, as in 
the salts NaHS and NaHS0 4 . These are called acid salts, since 






SULFUR AND ITS COMPOUNDS 195 

they are at once both salts and acids. Acid salts are often 
designated by the prefix bi- ; thus, NaHS0 4 is called sodium 
bisulfate, sodium acid sulfate, or sodium hydrogen sulfate. 

Carbon disulfide (CS 2 ). When sulfur vapor is passed over 
highly heated carbon the two elements combine, forming 
carbon disulfide (CS ), just as oxygen and carbon unite to 
form carbon dioxide (C0 2 ). The substance is a heavy, color- 
less liquid possessing, when pure, a pleasant odor. On stand- 
ing, especially when exposed to sunlight, it undergoes a slight 
decomposition and acquires a most disagreeable odor. It is a 
very good solvent for many substances, such as gums, rubber, 
waxes, and fats, which are insoluble in most liquids. It boils at 
a low temperature (46°), and its vapor is not only very poi- 
sonous but is very inflammable. It is prepared in considerable 
quantities for use as a solvent and as an insecticide. 

EXERCISES 

1. What elements so far studied occur in different allotropic forms? 

2. Suppose you were to grind a part of a stick of brimstone to 
a fine powder. Would the powder and the stick be allotropic forms 
of sulfur? 

3. How could you tell whether any two forms of sulfur are allotropic 
forms or simply different physical states of the same form ? 

4. Why were the ancients so familiar with sulfur? 

5. The reaction used in preparing hydrogen sulfide in the laboratory 
is a reversible one. Why does it ordinarily complete itself ? 

6. If hydrogen sulfide were a liquid, could we use the ordinary 
method of preparation? 

7. Hydrogen sulfide is said to be as poisonous as carbon monoxide, 
and yet we do not fear it as we do carbon monoxide. Why? 

8. Do sulfur waters contain free sulfur ? 

9. (a) What foods contain sulfur? (b) Why do such foods blacken 
silver spoons in contact with them ? 

10. Enumerate the different reactions so far studied in which catalytic 
agents have been employed. 



196 CHEMISTRY AND ITS USES 

11. Can you suggest any reason why sulfur dioxide is so much more 
soluble in water than is oxygen or nitrogen ? 

12. How should you expect dilute sulfuric acid to act (a) upon iron? 
(b) upon silver ? (See displacement series.) 

13. Zinc dissolves in both dilute and concentrated sulfuric acid. 
(a) What compound of the metal forms in each case ? (b) What gas is 
evolved in each case ? 

14. What reaction should you expect to take place if hydrogen sulfide 
were passed into a solution of sodium hydroxide ? 

15. (a) Define each of the following terms : acid, salt, monobasic 
acid, tribasic acid, normal salt, acid salt, (b) Give an example of each. 

16. State what compound a metal would form when dissolved in each 
of the following : hydrochloric acid, nitric acid, sulfuric acid, aqua regia. 

17. When carbon disulfide burns, the two elements present form the 
oxides C0 2 and S0 2 . Write the equation for the reaction. 

18. (a) The compound Na 2 S0 3 is a salt of what acid? (6) When 
this salt is brought in contact with sulfuric acid, sulfur dioxide is 
evolved ; account for its formation. 

19. The common powders sold for cleaning straw hats contain 
sodium sulfite. Upon what property does its cleaning action depend ? 

20. What volume of oxygen would be required to burn completely 
100 liters of hydrogen sulfide ? (Note that both hydrogen sulfide and 
oxygen are gases. See p. 115.) 

21. Calculate the percentage composition of hydrogen sulfate. 

22. During the World War one plant produced daily 1000 tons of 
sulfuric acid containing 60 per cent of hydrogen sulfate. What was the 
daily consumption of sulfur in this plant, assuming that none was lost 
in the process ? 

23. What weight of carbon disulfide can be made from 100 kilograms 
of sulfur? 



CHAPTER XXI 
THE PERIODIC LAW 

A number of the elements have now been studied some- 
what closely. Of these oxygen, hydrogen, nitrogen, carbon, 
chlorine, and sodium have almost no point of similarity as 
regards their chemical conduct. On the other hand, oxygen 
and sulfur, while quite different physically, have much in 
common in the way they act toward other chemicals. 

More than eighty elements are now known. If all these 
should have properties as diverse as do oxygen, hydrogen, 
nitrogen, carbon, chlorine, and sodium, the study of chemistry 
would plainly be very difficult and complicated. Fortunately 
a study of the elements shows that certain ones resemble each 
other more or less closely, so that it is possible to arrange them 
in groups and then study the group as a whole. A number of 
different methods for classifying the elements have been sug- 
gested, but the one advanced in 1869 by the Russian chemist 
Mendeleeff (Fig. Ill) has proved the most fruitful. In accord- 
ance with this method the elements are arranged in groups or 
periods, according to their atomic weights. 

The periodic grouping. The general arrangement suggested 
by Mendeleeff and extended so as to include elements more 
recently discovered is as follows : Omitting hydrogen, which 
has the smallest atomic weight, and beginning with helium, 
which has an atomic weight of 4.00, the succeeding seven 
elements are arranged in a horizontal row in the order of 
their atomic weights, as given below. These eight elements 
all differ markedly from each other, but neon, the one hav- 
ing the next highest atomic weight, is very similar to helium. 

197 



198 



CHEMISTRY AND ITS USES 



It is placed just under helium, and a new horizontal row fol- 
lows, as shown below. The element following chlorine, namely, 
argon, resembles helium and neon and begins a third row. 

He (4.00) Li (6.94) Gl (9.1) B (10.9) C (12.005) N (14.008) O (16) F (19) 
Ne (20.2) Na (23) Mg (24.32) Al (27.1) Si (28.3) . P (31.04) S (32.06) CI (35.46) 
A (39.90) K(39.1) Ca (40.07) Sc (45.1) Ti (48.1) V (51) Cr (52) Mn (54.93) 

If now we consider the elements that fall under each other 
in these three rows, a remarkable fact is brought to light. 

Not only is there a strong simi- 
larity between helium, neon, and 
argon, which form the first ver- 
tical column, but the elements in 
the other columns exhibit much 
of the same hind of similarity 
among themselves and evidently 
form natural groups. 

Iron, nickel, and cobalt, fol- 
lowing manganese, have atomic 
weights near together and are 
very similar chemically. They 
do not strongly resemble any of 
the elements so far considered 
and so are placed in a group 
by themselves. The first three 
horizontal rows of the table 
(p. 199) show the arrangement 
of these twenty-seven elements. A new horizontal row is begun 
with copper. Following the fifth and seventh rows are groups 
of three closely related elements, so the completed arrange- 
ment has the appearance represented in the table. 

Relation of properties of elements to atomic weights. There is 
evidently a close relation between the properties of an element 
and its atomic weight. For example, consider the elements in 
the first horizontal row. Helium is an inert element. Following 




Fig. 111. Mendeteeff (1834-1907) 

A Russian chemist who proposed the 
periodic classification of the elements 



THE PERIODIC LAW 



199 



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200 CHEMISTEY AND ITS USES 

it, lithium is a metallic element, has a valence of 1, and possesses 
a strong base-forming character. The next element, glucinum, 
has a valence of 2 and is less strongly base-forming, while boron 
has some base-forming and some acid-forming properties. In 
carbon all base-forming properties have disappeared, and the 
acid-forming properties are more marked than in boron. These 
become still more emphasized as we pass through nitrogen and 
oxygen, until, on reaching fluorine, we have one of the strongest 
acid-forming elements. The properties of these eight elements 
vary regularly with their atomic weights or, in mathematical 
language, are regular functions of them. 

The periodic law. If it were true that helium had the 
smallest atomic weight of any of the elements and fluorine the 
greatest, so that in passing from one to the other we included 
all the elements, we could say that the properties of elements 
were regular functions of their atomic weights. But fluorine 
is an element of relatively small atomic weight, and the one 
following it, neon, breaks the regular order, for in it reappear 
all the characteristic properties of helium. Sodium, following 
neon, bears much the same relation to lithium that neon does 
to helium, and the properties of the elements in the second 
row vary much as the properties of the elements in the first 
row did, until argon is reached, when another repetition begins. 
The properties of the elements do not vary continuously, there- 
fore, with atomic weights, but at regular intervals there is a 
repetition, or period. This generalization is known as the peri- 
odic law and may be stated thus : The properties of elements are 
'periodic functions of their atomic weights. 

Two families in a group. The elements of each group 
(excepting Group 0) fall naturally into two families. The 
elements in the odd-numbered horizontal rows, or periods, 
form one family, those in the even-numbered periods the other. 
In the table these are arranged under the headings A and B. 
The elements in one family are much more similar to each 



THE PERIODIC LAW 201 

other than they are to those in the other family in the same 
group. Thus, magnesium, zinc, cadmium, and mercury form one 
family of very similar elements in Group II, while glucinum, 
calcium, strontium, barium, and radium form the other. 

Family resemblances. Let us inquire more closely in what 
respects the elements of a family resemble each other. 

1. Valence. In general the highest valence of the elements 
in a family is the same, and the formulas of their compounds 
are therefore similar. The formulas R 9 0, RO, etc., placed 
below the columns, represent the oxides of the elements in 
the column, while the formulas RH, RH 2 , etc. represent the 
hydrides or chlorides, in which R represents any one of the 
elements of the family. 

2. Chemical conduct. The chemical characteristics of the 
members of a family are somewhat similar. If one member is 
a metal, the others usually are ; if one is a nonmetal, so, too, 
are the others. There is also a certain regularity in the 
conduct of the elements in each family. 

3. Physical properties. In the same way, the physical prop- 
erties of the members of a family are in general somewhat 
similar and show a regular gradation as we pass from element 
to element in the family. 

Value of the periodic law. The periodic law has proved of 
much value in the development of the science of chemistry. 

1. It simplifies study. It is at once evident that such regu- 
larities very much simplify the study of chemistry. A thorough 
study of one element of a family makes the study of the other 
members a much easier task. 

2. It suggests the probable existence of new elements. When 
the periodic law was first formulated there were a number of 
vacant places in the table which evidently belonged to ele- 
ments at that time unknown. From their position in the table 
Mendeleeff predicted with great precision the properties of the 
elements which he felt sure would one day be discovered to 



202 CHEMISTRY AND ITS USES 

fill these places. Three of them, scandium, germanium, and 
gallium, were found within fifteen years, and their properties 
agreed in a remarkable way with the predictions of MendeleefT. 

3. It indicates probable errors. The physical constants of 
many of the elements did not at first agree with those sug- 
gested by the periodic law, and a further study of many such 
cases showed that errors had been made. 

Imperfections of the law. There still remain a good many 
features which must be regarded as imperfections in the law. 
Most conspicuous is the fact that the element hydrogen has 
no place in the table. Moreover, according to their atomic 
weights, tellurium should follow iodine, and argon should 
follow potassium, but their properties show hi each case that 
this order must be reversed. 

The periodic law is therefore to be regarded as but a partial 
and imperfect expression of some very important and funda- 
mental relation between the substances which we know as 
elements. It has been proved in recent years that the atoms 
of the elements are complex bodies made up of common units, 
the number of units varying with the element. This conception 
throws much light upon the periodic law, but the discussion 
of the subject is beyond the scope of this text. 

EXERCISES 

1. Can you suggest any occurrences in nature that are periodic in 
character ? 

2. In what respect are the elements in Group (see page 199) alike? 

3. Chlorine, bromine, and iodine occur in Group VII. (a) Having 
studied chlorine, what compounds should you expect bromine and iodine 
to form with hydrogen ? (b) What should you expect the nature of these 
compounds to be ? 

4. Suppose an element were discovered having an atomic weight of 
about 60. Where would it find a place in the table ? What properties 
would it probably have ? 







CHAPTER XXII 






THE CHLORINE FAMILY 

• 




Name 


Atomic 
Weight 


Melting 
Point 


Boiling 
Point 


Color and State 


Fluorine (F) 




19.00 


— 223° 


-187° 


Pale-yellowish gas 


Chlorine (CI) . 




35.46 


-101.5° 


-33.6° 


Greenish-yellow gas 


Bromine (Br) . 




79.92 


-7.3° 


63° 


Red liquid 


Iodine (I) . . 




126.92 


113.5° 


184.4° 


Purplish-black solid 



The family. The four elements named in the above table 
form a strongly marked family and illustrate very clearly the 
way in which the members of a family in a periodic group 
resemble each other, as well as the character of the differences 
which we may expect to find between the individual members. 
The elements constituting the family are often termed the 
halogens, a word meaning " producers of salt." Chlorine has 
already been discussed in Chapter XV. The student should 
review that chapter in connection with the present one. 



Fluorine 

General discussion. It is safe to say that the student of this 
text has never seen fluorine, and, indeed, very few persons have. 
This is due to the fact that the element is difficult to prepare 
and has no uses. Nevertheless its compounds are of consider- 
able importance. Fluorine is a pale-yellowish gas. It is ex- 
tremely active, even at ordinary temperatures, and is very 
poisonous. It combines with all the common elements save 
oxygen. It has a very strong affinity for hydrogen, decom- 
posing water with violence and combining with hydrogen to 

203 



204 



CHEMISTRY AND ITS USES 



form a gas known as hydrogen fluoride (H 2 F 2 ). Of course 
such an active element would not occur free in nature. 
Its most important natural compounds are the minerals 
fluorite (CaF 2 ) and cryolite (Na 3 AlF 6 ). Traces of fluorides 
are also to be found in bones and in the enamel of the teeth. 




Fig. 112. Tablet erected by the associates and friends of Moissan in 

his laboratory in Paris in 1906, on the twentieth anniversary of the 

isolation of fluorine 

Preparation. Because of its great activity fluorine is diffi- 
cult to liberate from its compounds. The French chemist 
Moissan (Fig. 112) in 1886 finally succeeded in obtaining it 
in a free state. 

Moissan obtained the element by electrolyzing hydrogen fluor- 
ide at low temperatures. Pure hydrogen fluoride is a noncon- 
ductor; hence Moissan dissolved in it a little potassium acid 
fluoride (KHF 2 ) to render it a conductor. A much simpler method 
of preparation consists in electrolyzing melted potassium acid 
fluoride (KHF 2 ). Moissan carried out his experiments in a vessel 
made of platinum — a metal which is very inactive but very 
expensive. Copper vessels, however, will do nearly as well and 
are much cheaper. 



THE CHLORINE FAMILY 



205 



Hydrogen fluoride (H 2 F 2 ) ; hydrofluoric acid. Hydrogen fluor- 
ide is a low-boiling liquid (boiling point, 19.4°). It is prepared 
by the action of sulfuric acid on the mineral fluorite (CaF 2 ), 
thus: 



CaF„ 



H 2 S0 4 



CaS0 4 + H 2 F 2 



It is soluble in all proportions in water, forming hydrofluoric 
acid. This is a strong acid and is extremely corrosive. It 
readily reacts with sand and powdered 
quartz, both of which consist of silicon 
dioxide (Si0 2 ), to form water and the 
gas silicon tetrafluoride (SiF 4 ), thus : 

> SiF, + 2 HO 




Si0 2 +2H 2 F 2 

This property makes possible the chief 
use of hydrofluoric acid ; namely, the 
etching of glass, all varieties of which 
contain silicon dioxide either free or 
in combination. The acid is preserved 
in bottles made of a wax obtained 
from petroleum and known as ceresin 
(Fig. 113). 

Fig. 113. A bottle made 
Etching of glass. The glass vessel to out of ceresin for ho iding 

be etched is painted over with a protec- hydrofluoric acid 

tive paint upon which the acid will not 

act, the parts which it is desired to make opaque being left unpro- 
tected. A mixture of fluorite and sulfuric acid is then painted 
over the vessel, and after a few minutes the vessel is washed clean. 
Wherever the hydrofluoric acid comes in contact with the glass, 
it acts upon it, destroying its luster and making it opaque, so that 
the exposed design will be etched upon the clear glass. Frosted- 
glass globes are often made in this way, but more frequently by 
a sand blast. The etching may also be effected by covering the 
glass with a thin layer of paraffin, cutting the design through the 
wax, and then exposing the glass to the fumes of the gas. 



206 



CHEMISTRY AND ITS USES 



Bromine 

General discussion. This element, discovered by the French 
chemist Ballard in 1826, is a dark-red heavy liquid which 
boils at 63° and freezes at — 7.3°. It vaporizes readily at ordi- 
nary temperatures. This vapor has an offensive odor and is 
very irritating to the eyes and throat. One volume of the 
liquid dissolves in 100 volumes of water. It resembles chlorine 
in its chemical conduct except that it is less active. Bromine 
occurs in combination with metals, sodium bromide (NaBr) 
and magnesium bromide (MgBr 2 ) being the most abundant 
forms. These are found in many salt waters. The salt waters 
obtained by sinking deep wells along the Ohio River and in 
certain parts of Michigan are the richest in bromine and con- 
stitute the source from which the element is prepared. 

Preparation. The methods used for preparing chlorine in the 
laboratory (p. 137) serve also for the preparation of bromine. 
Commercially it is chiefly prepared by passing an electric 

current through salt waters 
containing bromides. Bro- 
mine and some chlorine 
are liberated ; the chlorine, 
however, assists in the reac- 
tion, since it decomposes the 
unchanged bromides, liber- 
ating bromine, thus: 

NaBr + CI ^NaCl + Br 




Fig. 114. An apparatus for the prepa- 
ration of bromine in the laboratory 



Laboratory preparation of 
bromine. A mixture of sodium 
bromide and manganese diox- 
ide is introduced into the retort A (Fig. 114). Sulfuric acid is added 
and the mixture heated. Bromine is liberated and condenses in 
the flask B, kept cool by ice water in the beaker C. The equation is 

2 NaBr + 2 H 2 S0 4 + Mn0 2 >■ Na 2 S0 4 + MnS0 4 + 2 H 2 + Br 2 



THE CHLORINE FAMILY 207 

Hydrogen bromide (HBr) ; hydrobromic acid. Hydrogen 
bromide is a colorless gas resembling hydrogen fluoride and 
hydrogen chloride except that it is less stable (that is, more 
easily decomposed). Its solution in water constitutes hydro- 
bromic acid. The salts of hydrobromic acid are known as 
bromides. They resemble the chlorides in their properties. 

Uses. Bromine is used in making certain dyes and medi- 
cines and in the preparation of bromides. Silver bromide 
(AgBr) is used extensively in photography, while sodium 
bromide and potassium bromide are common drugs. 

Bromine in the World War. Among the poison gases used in the 
World War were certain compounds known as lachrymators be- 
cause they caused inflammation of the eyes and a copious flow of 
tears. Bromine is the essential constituent of the most effective 
of these lachrymators. To obtain supplies of bromine necessary 
for their production the United States, during the war, sunk 
sixteen wells in the vicinity of Midland, Michigan. These wells 
represent a total possible output of bromine amounting to about 
500,000 pounds annually. 

Iodine 

General discussion. Iodine was discovered in the ashes of 
sea plants by the French chemist Courtois in 1812. It differs 
from the other halogens in that it is a purplish-black, shining 
crystalline solid. When warmed it gives off a beautiful violet 
vapor. It is only very slightly soluble in water, but readily 
dissolves in carbon tetrachloride and in alcohol. It resembles 
chlorine in its chemical conduct, but is even less active than 
bromine. Even a very small amount of free iodine colors 
starch blue — a reaction which is used to detect the presence 
of both iodine and starch. Small amounts of iodine are found 
combined with metals in sea water, from which it is absorbed 
by certain sea plants, so that it is found in their ashes. It 
is also found in Chile as a constituent of sodium nitrate 
(Chile saltpeter). 



208 . CHEMISTRY AND ITS USES 

Preparation. In the laboratory iodine is prepared in the same 
way as bromine, using Nal in place of NaBr (p. 206). Com- 
mercially iodine was formerly obtained entirely from the ashes 
of seaweed (kelp), and a limited quantity is still obtained from 
this source. Most of our supply comes from the beds of Chile 
saltpeter. 

Hydrogen iodide (HI); hydriodic acid. Hydrogen iodide 
is a colorless gas, but is less stable than the correspond- 
ing compounds of the other halogens. Its solution in water 
constitutes hydriodic acid. The salts of this acid are called 
iodides, and these in general resemble the chlorides and 
bromides. 

Use. A solution of iodine in alcohol is known as tincture 
of iodine and constitutes one of our most common antiseptics. 
Potassium iodide (KI) is used in medicine, and iodine has 
many uses in chemical analysis. 

Properties of the halogens contrasted. By studying the table 
at the head of this chapter it will be noted that the halogens 
exhibit a regular gradation of properties. Fluorine is a light, 
slightly colored gas, chlorine is a heavier gas with more marked 
color, while bromine is a red liquid, and iodine a nearly black 
solid. Their melting points and boiling points increase with their 
atomic weights (see table). Their affinity for hydrogen and the 
metals, on the other hand, decreases with their atomic weights, 
that of fluorine being the greatest and iodine the least. Fluorine 
and bromine do not combine with oxygen, while chlorine and 
iodine form unstable oxides. 

Gay-Lussac's Law of Volumes 

Experiments show that 1 volume of hydrogen combines with 
1 volume of chlorine to form 2 volumes of hydrogen chloride. Similar 
combining ratios hold good when hydrogen combines with the 
vapor of bromine and of iodine. These facts recall the simple vol- 
ume relations already noted in the study of the composition of 



THE CHLORINE FAMILY 



209 



steam (p. 71). These relations may be represented graphically 
hi the following way, the squares representing equal volumes: 





H 2 


+ 


Cl 2 




TTP1 


j_ urn 




















H 2 


| H 2 


+ 


o 2 




H 2 


+ H 2 





In the early part of the past century the distinguished 
French chemist Gay-Lussac (Fig. 26) studied the volume 
relations of many combining gases and concluded that simi- 
lar relations always hold. His observations are summed up 
in the following generalization, known as the law of volumes : 
When two gases combine chemically there is always a simple ratio 
between the volumes that combine and also between the volume of 
either one of them and that of the product, provided it is a gas. 
By a simple ratio is meant, of course, the ratio of integer 
numbers ; as, 1 : 2 or 2 : 3. 



EXERCISES 

1. Give the name and the nationality of the discoverer of each of 
the halogens. 

2. Contrast the properties of the halogens. 

3. Contrast the chemical conduct of the halogens. 

4. (a) Give the names and formulas of the compounds that each of 
the halogens forms with hydrogen, (b) To what class of compounds do 
they belong ? 

5. What elements are liquids at ordinary temperature? 

6. Consult the dictionary for the significance of the names of each 
of the halogens. 

7. (a) How do you account for the fact that the pure liquid hydro- 
gen fluoride is a nonconductor of electricity? (&) How did Moissan 
render it a conductor ? 

8. In what other connection has the name of Moissan been 
mentioned ? 

9. Why cannot fluorine be prepared by electrolyzing hydrofluoric 
acid? 



210 CHEMISTRY AND ITS USES 

10. Why do we write the formula for hydrogen fluoride as H 2 F 2 
while that for hydrogen chloride is written HC1 ? 

11. (a) What gas has been studied that resembles the vapor of 
bromine in color ? (b) How could you distinguish between the two ? 

12. Why do solutions of hydrogen bromide and hydrogen iodide color 
on standing, while hydrogen fluoride and hydrogen chloride do not ? 

13. Sodium chloride, sodium bromide, and sodium iodide are all 
white crystalline solids. How could you distinguish between them ? 

14. What reaction would take place if chlorine were passed into a 
solution («) of sodium bromide ? (&) of sodium iodide ? 

15. Suppose that iodine were added to a solution of sodium chloride. 
Would any reaction take place ? 

16. Enumerate the acids so far studied, giving the composition 
of each. 

17. Two volumes of hydrogen combine with one volume of oxygen 
to form two volumes of water vapor (see Gay-Lussac's law of volumes, 
p. 209). Show that this fact is in accord with Avogadro's principle (p. 47). 

18. A certain mineral is known to be either rock salt (NaCl) or 
fluorite. How could you decide which it is ? 

19. What weight of fluorite is necessary for preparing 100 g. of 
hydrofluoric acid containing 50 per cent of hydrogen fluoride? 

20. The salt water used in preparing bromine contains 0.12 per cent of 
bromine ; 1 ton of this brine would yield how many pounds of bromine ? 

21. What weight of sodium iodide would be required for the 
preparation of 10 g. of iodine? 



CHAPTER XXIII 
HYDROCARBONS ; PETROLEUM 

Hydrocarbons. Carbon and hydrogen unite to form a great 
many compounds, and these are known as the hydrocarbons. 
Their importance may be inferred from the fact that, mixed 
together in varying proportions, they constitute such valu- 
able substances as natural gas, gasoline, kerosene, lubricating oils, 
vaseline, and paraffin. 

In order to simplify the study of the hydrocarbons it is 
convenient to arrange them in groups, or series. The most 
extensive of these is the methane series. In the following 
table are given the names, formulas, and boiling points of 
some of the members of this series : 





Boiling 




Boiling 




Point 




Point 


Methane (CH 4 ) . . 


. - 160° 


Pentane (C 5 H 12 ) 


+ 36° 


Ethane (C 2 H 6 ) . . 


. -93° 


Hexane (C 6 H 14 ) 


+ 69° 


Propane (C 3 H 8 ) . . 


. -45° 


Heptane (C 7 H 16 ) 


+ 98° 


Butane (C 4 H 10 ) . . 


+ 1° 


Octane (C 8 H 18 ) . . 


. +125.5° 



General formula (C n H 2n + 2 ) 

It will be noted that each member of this series differs from 
the one preceding it by the group of atoms CH 2 , and that 
the boiling points gradually increase. All the members of this 
series are known up to the one having the formula C 28 H 5S . 
The lower members are colorless gases, the intermediate mem- 
bers are liquids, and the higher members are solids. They 
are insoluble in water and are all combustible, the carbon 
present burning to carbon dioxide and the hydrogen to water. 

211 



212 



CHEMISTRY AND ITS USES 



Petroleum and products derived from it. Petroleum is a 
dark-colored liquid found in the earth in certain localities. 
This liquid is composed principally of liquid hydrocarbons in 
which are dissolved both gaseous and solid hydrocarbons. 
Crude petroleum is not only used as fuel in this country, par- 
ticularly on locomotives and ships, but many useful products 
are obtained from it by the process of refining, among them gaso- 
line, kerosene, lubricating oils, vaseline, and paraffin (Fig. 117). 
These products are not single compounds, but are mixtures of 
compounds boiling between certain limits. When pure they 
are all colorless. Sometimes the commercial products are 
colored by impurities. 



The refining of petroleum. In this process the crude oil is run 
into large iron stills (Fig. 116) and is then distilled. The dis- 
tillates which pass over between certain limits of temperature are 
kept separate and, after being purified by treatment with sulfuric 
acid and then with sodium hydroxide and water, serve for different 
uses. The liquid passing over between approximately 70° and 150° 
is known &snaphtha, while that passing over between 150° and 300° 
constitutes ordinary kerosene. A number of different naphthas are 
recognized commercially, differing in boiling points and density. 
Those of low boiling-point constitute ordinary gasoline and are 
used as a fuel in stoves and motors ; those of higher boiling- 
points are used in making paints. Benzine is a low-boiling naphtha 
and, being a good solvent for such organic substances as fats and 
oils, is used in cleaning fabrics (dry-cleaning). 

The liquid remaining after the kerosene and higher-boiling oils 
have distilled over is chilled, whereupon the solid constituents dis- 
solved in the oil separate. These are filtered off and constitute 
ordinary paraffin. The filtrate is then distilled, and from it various 
lubricating oils are obtained. 

Formerly kerosene was the most important of the products 
obtained from petroleum. At present, however, gasoline is the 
most in demand, so that every effort is made to increase the 
yield. To accomplish this efficiently the distillation is carried on 
under conditions that tend to decompose the heavier molecules 




Fig. 115. Typical scene in an oil field 

Wells are sunk into the oil-bearing strata, which vary in depth from a few 
feet to 4000 feet or more. Sometimes the oil in the strata is under such 
pressure that it will flow from the well in large volumes under its own 
pressure ; more often the oil must be pumped to the earth's surface. This 
crude oil, commonly called petroleum, is stored in tanks (see figure) until 
it is desired to refine it. In the United States the chief oil-producing 
regions at present are Texas, California, Oklahoma, and Pennsylvania. 
The total production varies from year to year, but is gradually increasing 
with the greater demand for gasoline and other products obtained from 
the oil. At present between 4,000,000 and 5,000,000 barrels of petroleum 
are produced annually, while the annual gasoline production amounts 
to nearly 4,000,000,000 gallons 




Fig. 116. View of stills for refining petroleum 

Sometimes the petroleum is used directly as a fuel. More often it is refined. 
This process consists in separating it into different constituents by distil- 
lation and purifying these products chiefly by washing them first with 
sulfuric acid, then with sodium hydroxide, and finally with water. The 
distillation is carried on in large iron stills, shown in the figure 








Fig. 117. Crude oil and the chief products obtained from it 



HYDROCARBONS; PETROLEUM 213 

making up the higher-boiling liquids into the simpler molecules 
which constitute liquids of lower boiling-points. The process is 
known as the cracking of oils. 

Use of benzine and gasoline in our homes. Because of the ease 
with which benzine and gasoline (these two terms are often used 
interchangeably) burn, and also because of the explosive character 
of their vapor when mixed with air, accidents often result from 
their use, especially when they are employed in our, homes for 
cleaning fabrics. Recent statistics show that in a single year 
1040 persons were burned to death and 3120 injured as a result 
of gasoline fires. It is obvious, therefore, that the greatest care 
should be exercised in its use. 

Fractional distillation. The process of purifying a mixture 
such as petroleum is known as fractional distillation. In this 
process the mixture is heated, and in a general way the com- 
pounds distill over in the order of their boiling points. By re- 
peating the process it is ordinarily possible to separate a mixture 
into the compounds present. In the refining of petroleum it is 
not necessary to effect such a sharp separation, but only to 
obtain a mixture of compounds which boil within certain limits. 

Asphalt. When petroleum from certain sources is distilled 
there is left in the retort a heavy, tarry liquid. A similar sub- 
stance is found in nature, both in the liquid and the solid form, 
and is known as asphalt The most important deposit of asphalt 
is on the island of Trinidad (Fig. 118). Asphalt is used for 
road construction, for making paints, and for roofing. 

Methane (CH 4 ). Pure methane is a colorless, odorless gas 
about one half as heavy as air. It is commonly known as marsh 
gas, since it is formed in marshes and, in general, wherever 
organic matter decays or is heated in the absence of air. It 
constitutes about 30 per cent of coal gas and from 90 per cent 
to 95 per cent of natural gas. It often collects in mines, and 
when mixed with air is called fire damp. Such mixtures are 
very explosive. When ignited it burns with a pale-blue flame : 
CH 4 + 2 2 — >C0 2 + 2H 2 



214 



CHEMISTKY AND ITS USES 



Safety lamp. Fortunately the ignition point of fire damp (that 
is, the temperature at which it takes fire) is high and its flame 
may be extinguished by cooling. In 1815 Sir Humphry Davy 
(Fig. 86) invented a miner's lamp based on this principle, in which 
the usual chimney of a lantern is replaced by a wire gauze (Fig. 119). 
An explosion flame starting at the wick is so cooled by the metal 
wire that ignition ceases and the explosion is confined to the 




Fig. 118. Mining asphalt on Trinidad island for use in road construction 

interior of the lamp. The principle may be demonstrated by hold- 
ing a wire gauze a few inches above a Bunsen flame and parallel 
with the table (Fig. 120). When the gas is turned on and a light 
applied above the gauze, the resulting flame rests upon the gauze, 
but does not pass through it to the burner. 

Halogen derivatives of methane. As a rule the hydrogen present 
in a hydrocarbon may be displaced by a halogen element, atom 
for atom. In this way there are formed from methane a number 
of derivatives, the most important of which are the following : 

Chloroform (CHC1 3 ), a heavy, colorless liquid boiling at 61°, is 
the well-known ansesthetic used in surgery. Carbon tetrachloride 



HYDROCARBONS ; PETROLEUM 



215 



(CC1 4 ) resembles chloroform in appearance. It is a good solvent, 
especially for fatty substances. It is often used to remove grease 
spots from fabrics and is sold for this purpose under the name of car- 
bona. It possesses the advantage over benzine 
of being noninflammable, but it is more expen- 
sive. Large quantities are used as a fire extin- 
guisher (jpyrene) (p. 114). Iodoform (CHI 3 ) is 
a yellow crystalline solid and is largely used 
as an antiseptic in the treatment of wounds. 

Acetylene (C 2 H 2 ). In addition to the hy- 
drocarbons composing the methane series, 
many others are known. Among these is 
acetylene. This is a colorless gas having, 
when pure, a faint, pleasant odor. It is 
easily obtained by the action of water on 
calcium carbide (CaC 2 ) : 

CaC 2 + 2H 2 0-> C 2 H 2 + Ca(OH) 2 

In this way the gas is prepared in large quantities for use as 
an illuminant and as a source of intense heat. When heated 
it decomposes, with evolution of a great 
deal of heat: 




Fig. 119. The miner 
safety lamp 




C 2 H 2 — y 2 C + H 2 + 49,300 cal. 
By this method the Germans prepared 
much of the hydrogen used for filling 
the Zeppelins in the World War. 

When compressed in cylinders acety- 
lene is very explosive, since the heat 
liberated in compressing the gas is suffi- 
cient to start decomposition. With the 
proper admixture of air it burns with a 
brilliant white light. The flame is very hot, because to the heat 
of combustion of the hydrogen and the carbon there is added the 
heat of decomposition of the acetylene undergoing combustion : 
2 C 2 H 2 + 5 2 ►■ 4 C0 2 + 2 H 2 + 603,260 cal. 



Fig. 120. An experiment 
to illustrate the principle 
* of the safety lamp 



216 



CHEMISTRY AND ITS USES 



Uses of acetylene. Acetylene is chiefly used as an illumi- 
nant when electric lights are not available. It may be safely 
stored in metal cylinders by filling the cylinder with some 
porous material (such as asbestos and cotton), partially satu- 
rating this with a liquid compound called acetone, and then 



■1 






. 






1 1 


1 






%JS&^ 






I 


1 1 


1 




I 




| .;■ J 










#£- 


* ; ' ■■■"' s 














. ,' : \. 


w% m 








: 


I ^ 


^ 



Fig. 121. Cutting an iron plate by means of the oxyacetylene blowpipe 



forcing in the gas at low temperatures. Under pressure the 
acetone dissolves a large volume of the gas. In this form it 
is now a common article of commerce. 

The intense heat generated by the combustion of acetylene 
makes it useful in certain processes requiring high tempera- 
tures, such as burning the carbon out of motor engines and 
the welding and cutting of metals. For this purpose the acety- 
lene is burned in an apparatus known as the oxyacetylene 
blowpipe, which is very much like the blast lamp in principle. 
A temperature of about 2700° may be obtained in this way. 
This blowpipe has been found especially useful in cutting iron 
structures (Fig. 121), since the tip of the flame, when drawn 






HYDROCARBONS; PETROLEUM 217 

slowly over the metal, melts and burns it at the point of con- 
tact and thus makes it possible to cut the metal into pieces. 

Ethylene (C 2 H 4 ). Small percentages of this gas are present in 
coal gas. It is prepared by heating the vapor of ordinary alcohol 
(C 2 H 6 0) in contact with kaolin (a compound of aluminium), which 
acts as a catalytic agent : 

C 2 H 6 0— »C 2 H 4 + H 2 

During the World War the United States prepared enormous 
quantities of ethylene for use in the manufacture of the poison 
liquid known as mustard gas. This compound, while called a gas 
during the war, is really a high-boiling, nearly colorless liquid. 
It has the formula (C 2 H 4 C1) 2 S, and its chemical name is dichlor- 
ethyl sulfide. It is absorbed by the skin, and after absorption it 
decomposes, forming hydrogen chloride, which produces large 
blisters and causes severe and often fatal inflammation. The 
compound is made according to the following equation : 

2 C 2 H 4 + S 2 C1 — > (C 2 H 4 C1) 2 S + S 

EXERCISES 

1. Can the flame of gasoline be extinguished by water? 

2. How could you distinguish between gasoline and kerosene ? 

3. Why not burn gasoline in lamps? 

4. Why not substitute coal oil for gasoline in motor-car engines? 

5. Distinguish between fractional distillation and destructive dis- 
tillation. 

6. Methane is formed when organic matter decays under water, 
(a) Might this account for the formation of natural gas ? (b) What is 
the function of the water ? 

7. How do you account for the formation of carbon monoxide in a 
gasoline engine (p. 114) ? 

8. It is stated that in the gasoline engine, as ordinarily operated, 
about 30 per cent of the effective power of the gasoline is not recovered. 
Suggest a reason for this. 

9. Under what conditions is natural gas explosive ? 

10. Write the equations for the combustion of methane ; of acetylene. 



218 CHEMISTRY AND ITS USES 

11. In what proportions by volume ought acetylene and oxygen to 
be mixed in an oxyacetylene blowpipe to get the maximum heat? 

12. In a general way, to what extent is modern civilization dependent 
on petroleum? 

13. What products of petroleum are used in your home? 

14. Under what conditions would natural gas become a liquid? 

15. Kerosene is a good solvent for fats and oils. Why not use it in 
place of gasoline for cleaning fabrics? 

16. Why is it that kerosene is so much safer to handle than gasoline ? 

17. Suppose you were given two bottles, the one filled with hydrogen 
and the other with marsh gas, how would you proceed to determine 
which of the two contained the hydrogen ? 

18. Assuming that natural gas is pure methane, (a) what volume 
of oxygen would be required for the combustion of 1000 cu. ft. of the 
gas? (&) what volume of carbon dioxide would be produced? (Solve 
by inspection of equation, p. 115 .) 

19. (a) Write the equation for the combustion of octane. (&) Assum- 
ing that gasoline is pure octane and that 1 gallon of it weighs 3500 g., 
what weight of oxygen would be required for the complete combustion 
of 1 gallon of gasoline ? (c) What volume of air would be required to 
furnish this amount of oxygen ? 



CHAPTER XXIV 
FUELS; ELECTRIC FURNACES; FLAMES 

Fuels. A variety of substances are used as sources of heat, 
the most important of them being the various fuel gases, to- 
gether with coal, wood, and petroleum. The composition of 
a number of these fuel gases is given in the table on page 224. 
Most of them serve as illuminants as well as fuels. 

Coal gas. It has been known for several centuries that when 
soft (bituminous) coal is heated out of contact with air, com- 
bustible gases are formed ; indeed, gas obtained in this way 
was used for street lighting in London and Paris more than 
a hundred years ago. 

Manufacture of coal gas. The manufacture of coal gas is rep- 
resented in a diagrammatic way in Fig. 122. The coal is intro- 
duced into a closed retort A and heated by the fire below. A 
number of these retorts are placed in parallel rows, each being 
furnished with a delivery pipe. These pipes lead into a large 
pipe B (known as the hydraulic main), which runs at right angles 
to the retorts. When the coal is heated out of contact with air a 
large number of compounds are formed, some of which, at ordinary 
temperatures, are gases, others are liquids, and others are solids. 
These escape into the hydraulic main, and, the temperature being 
reduced somewhat, some of the products that are liquid or solid at 
ordinary temperature condense in the form of a black, tarry mass 
known as coal tar and are continuously drawn off. The gas then 
passes into the cooler C, where the remaining tar condenses. On 
the top of the tar there collects a liquid containing ammonia and 
known as ammoniac al liquor (p. 164). In the scrubber D the 
partially purified gas passes through a column of loose coke, over 
which water is sprayed, where it is still further cooled and to 

219 



220 



CHEMISTRY AND ITS USES 



some extent purified from soluble gases such as hydrogen sulfide 
and ammonia. In the purifier E it passes over a bed of lime or of 
iron oxide, which removes the remainder of the sulfur compounds, 
and from this it enters the large gas holder F, from which it is 
distributed to consumers. 

The great bulk of the carbon remains in the retort as coke and 
as retort carbon, The yield of gas, tar, and soluble materials 
depends upon many factors, such as the composition of the coal, 




Fig. 122. Diagram of a plant used for the manufacture of coal gas 
and its by-products 



the temperature employed, and the rate of heating. One ton of 
good gas coal yields approximately 10,000 cu. ft. of gas, 1400 lb. 
of coke, 120 lb. of tar, and 20 gal. of ammoniacal liquor. 

Not only is the ammonia obtained in the manufacture of the 
gas of great importance, but the coal tar is the source of many 
very useful substances, as will be explained later. 

The by-product coke oven. It will be observed that coke is formed 
in the process used in the manufacture of coal gas. Coke is a very 
important product and is used in large quantities, especially in the 
reduction of metals, such as iron, from their ores. The quantity 
of coke obtained in the manufacture of coal gas has never been 
sufficient to meet the demand. Much of the additional coke 
required has been prepared by coking the coal in ovens called 
beehive ovens because of their shape. The coking of coal in these 
ovens is carried out as follows : The oven is nearly filled with 
coal, and the coal is ignited. After the fire is well started the draft 
is shut off, and the heat formed in the combustion of a portion of 



FUELS; ELECTRIC FURNACES; FLAMES 221 

the coal is sufficient to coke the remainder of the coal. In this 
process all the coal tar, coal gas, and ammonia escape through an 
opening in the top of the furnace and are lost. 

The growing demand for ammonia, as well as for the products 
obtained from coal tar, has led to the construction of furnaces or 
ovens for the coking of coal which make it possible to save the 
coal tar and ammonia formed in the process. Such ovens are 




Fig. 123. View of a by-product coking plant 

The coal is heated in the narrow ovens. The view shows the coke being removed 
from one of the ovens 



known as by-product coke ovens, this term being chosen because 
the ammonia and coal tar formed in the process of coking the coal 
are by-products, the coke being the main product. Even at present 
more coke is made in the by-product ovens than in the beehive, 
and the percentage will undoubtedly increase since we cannot 
afford to waste such valuable products as are wasted in the 
beehive ovens. 

Fig. 123 represents a portion of a large by-product coking plant. 
It consists primarily of a number of narrow, upright ovens placed 
side by side, but separated sufficiently to admit of heating the coal 



222 CHEMISTRY AND ITS USES 

with which the ovens are filled. As a rule a portion of the gas 
generated in the process is used as a fuel for coking the coal. The 
hot flame of the burning gas strikes against the bottom and sides 
of the ovens. The coal tar and ammonia are separated and col- 
lected as in the manufacture of coal gas ; indeed, the by-product 
oven is a large coal-gas plant. When the process is complete 
each oven is pushed forward and the coke dropped into a car, as 
shown in the figure. 

Water gas. Water gas is essentially a mixture of carbon 
monoxide and hydrogen. It is made by passing steam over 
very hot anthracite coal or coke, when the reaction shown in 
the following equation takes place : 

C + H 2 ^CO + H 2 

The coal is burned in a draft of air until it is very hot. The 
air is then shut off and steam turned on. The temperature 
gradually falls because the reaction absorbs heat, and when it 
reaches about 1000° the steam is cut off and the air supply 
renewed. 

Water gas is very effective as a fuel, since both carbon mon- 
oxide and hydrogen burn with very hot flames. It has little 
odor and is very poisonous. Its use is therefore attended with 
some risk, since leaks in pipes are very likely to escape notice. 

Enriched water gas. When required merely for the production 
of heat, the gas as prepared above is at once ready for use. When 
made for illuminating purposes it must be enriched ; that is, illu- 
minants must be added, since both carbon monoxide and hydrogen 
burn with a nonluminous flame. This is accomplished by passing 
it into heaters containing highly heated petroleum oils. The gas 
takes up the gaseous hydrocarbons formed in the decomposition 
of the petroleum oils, and these hydrocarbons make it burn with 
a luminous flame. 

Producer gas. Producer gas is used in connection with many 
metallurgical furnace operations and also as a fuel for gas 
engines. It is made by burning coal under such conditions 






FUELS; ELECTRIC FUBNACES ; FLAMES 223 



that the product of combustion is largely carbon monoxide 
(Fig. 124). Very often a little steam is admitted with the air, 
and this on passing through the hot bed of coals is reduced as 
in the preparation of water gas. Made in this way, producer 
gas is composed mainly of 
carbon monoxide, hydro- 
gen, and nitrogen. It can 
be made from coal of a 
poor quality, even from 
lignite, and as gas engines 
run well with this gas, it 
furnishes the most econom- 
ical method for utilizing 
low-grade coal for power. 
Natural gas. In many 
regions of the United 
States, as well as in other 
countries, natural gas is 
obtained from wells drilled 
into a rock stratum hold- 
ing the gas. While it is 
variable in composition, it consists largely of methane, many 
samples containing as much as 95 per cent of this compound. 
It burns with a rather smoky name of moderate luminosity, 
but works well with a gas mantle. It has a high heat of 
combustion, as shown in the following equation : 

CH 4 + 2 2 >■ C0 2 + 2 H 2 4- 192,160 cal. 

It is an ideal fuel and is often conducted through pipes for 
hundreds of miles from the gas fields to cities. 

Comparative composition of fuel gases. The figures in the table 
on page 224 are the results of analyses of average samples, but 
since each kind of fuel gas varies considerably in composition, 
the values are to be taken as approximate only. 




Fig. 124. Diagram of the method of 
making producer gas 



224 



CHEMISTKY AND ITS USES 
COMPOSITION OF GASES 



Constituent 



H 2 

CH 4 

C 6 H 6 .... 
C 2 H 2 and C 2 H 4 . 

CO 

C0 2 

N 2 ..... 
2 ..... 
Other hydrocarbons 



Ohio 

Natural 

Gas 



0.9 
*9.5 
9.3 
0.3 
0.4 
0.3 
0.2 
0.0 
0.0 



Coal, 

GAS 



41.3 

43.6 

3.9 
6.4 
2.0 
1.2 
0.3 
1.5 



Water 
Gas 



52.88 
2.16 



36.80 
3.47 
4.69 



Enriched 

Water 

Gas 



37.96 
7.09 
2.01 
9.40 

32.25 
4.73 
3.96 
0.60 
1.80 



Producer 
Gas 



10.90 



0.60 
20.10 

8.50 
59.90 



Products of the combustion of ordinary fuels. Ordinary fuels, 
such as oil, wood, coal, and fuel gases, are largely made up of 

carbon and hydrogen, 
or their compounds. 
The chief products of 
the combustion of such 
fuels are carbon diox- 
ide and water. Asso- 
ciated with these are 
small amounts of other 
products, such as car- 
bon monoxide and sul- 
fur dioxide, the latter 
being formed from traces 
of sulfur compounds in 
the fuels. 

Rooms are not infre- 
quently heated by gas 
or oil stoves, with no provision for removing the products of 
combustion. Likewise, natural gas is often burned in stoves 
or grates with the damper closed so as to leave no opening into 




Fig. 125. Determining the calorific value of 

coal in the laboratory of the American Rolling 

Mill Company 



FUELS; ELECTRIC FURNACES; FLAMES 225 



the chimney. Such practices are greatly to be condemned, 
since the air in the rooms heated in this way soon becomes 
saturated with water vapor 
and so contaminated with 
the other products of com- 
bustion as to render it unfit 
for respiration. In walking 
along a street in cold weather 
one can pick out the rooms 
that are heated in this way, 
for the windows of such 
rooms are covered with a 
film of moisture. 

Calorific value of fuels. 
The various materials used 
as fuels differ much in the 
heat which they give out 
when burned. While many 
other factors are concerned 
in the value of a fuel, the 
chief one is its heat of com- 
bustion. The heat evolved by 
the combustion of one gram of a 
fuel is called its calorific value. 
In large contracts the price 
paid for a fuel is generally 
based on its calorific value, as 
well as upon its adaptability 

to the use to which it is to be put. The calorific value of (say) 
coal is determined by burning a definite weight of the sample 
in a bomb calorimeter constructed on the principle shown in 
Fig. 37, but fitted with a top so that it can be kept closed 
during the combustion (Fig. 126). The table on. page 226 
will give some average values for a few common fuels: 




Fig. 126. A bomb calorimeter for de- 
termining the calorific value of coal 

The coal is placed in the capsule C. The 
calorimeter is then filled with oxygen 
under pressure and the coal ignited by an 
electric current. After the coal is burned, 
the increase in the temperature of the 
water is noted and the calorific value 
deduced from this 



226 



CHEMISTRY AKD ITS USES 




CALORIFIC VALUE OF FUELS 

Wood (air-dried) about 3800-4000 cal. 

Bituminous coal (Pennsylvania), 35% volatile 

matter, 6 % ash about 8300 cal. 

Bituminous coal (Pocahontas), 18% volatile 

matter, 6% ash about 8700 cal. 

Anthracite coal (Connellsville), 12% ash . . . about 7300 cal. 

Coke, 10% ash about 7300 cal. 

The electric furnace. In recent years electric furnaces have come 
into wide use in operations requiring a very high temperature. 
Temperatures as high as 3500° can easily be reached. These 

furnaces are constructed on one of 
two general principles. 

1. Arc furnaces. In the one type 
the source of heat is an electric arc 
formed between carbon electrodes 
separated a little from each other, as 
shown in Eig. 127. The substance to 
be heated is placed in a vessel, usually 
a graphite crucible, just below the arc. 

2. Resistance furnaces. In the other 
type of furnace the heat is generated by the resistance offered to 
the current in its passage through the furnace. A typical form of 
such a furnace is illustrated in Fig. 60, which is used in the 
manufacture of graphite. 

Conditions necessary for flames. When one of the substances 
undergoing combustion remains solid at the temperature occa- 
sioned by the combustion, light may be given off, but there 
is no flame. Thus, iron wire burning in oxygen throws off 
a shower of sparks, but no flame is seen. When, however, 
both of the substances involved are gases or vapors at the 
temperature reached in the combustion, the act of union is 
accompanied by & flame. 

Flames from burning liquids or solids. Many substances which 
are liquids or solids at ordinary temperatures burn with a flame 
because the heat of combustion slowly vaporizes them, and the 



Fig. 127. An arc furnace 



FUELS; ELECTRIC FURNACES; FLAMES 227 



flame is due to the union of this vapor with the oxygen of the air. 
This may be shown in the case of a candle flame by holding one 
end of a slender glass tube in the base of the flame (Fig. 129). 
The unbuVned vapor in the inner part of the flame is thus con- 
ducted away and may be ignited at the upper end of the tube. 




Fig. 128. View of an electric furnace used by Moissan, who was one of the 
first scientists to experiment with reactions at high temperatures 

From Duncan's " Chemistry of Commerce " 

Structure of a flame. When hydrogen or carbon monoxide 
burns in oxygen, but one reaction takes place, and as a result 
the flame is very simple in structure. It consists of a colorless 
inner cone of unburned gas and an outer cone in which the 
union between the hydrogen and the oxygen is taking place. 
It follows that the outer cone is the hot part of the flame. That 
the inner cone is cool is shown by the fact that a match head 
suspended in this region (Fig. 130) before lighting the gas, 
and left there while the gas burns, is not ignited. 



228 



CHEMISTRY AND ITS USES 




Fig. 129. Method of proving that 

the interior of a candle flame 

contains combustible vapors 



The flames produced by the combustion of hydrocarbons 
such as are present in coal gas and natural gas or of mixtures 
of hydrocarbons with stearic acid, as in candles, is much more 

complex because several con- 
secutive reactions take place. 
For example, in the candle flame 
(Fig. 131) there are, broadly 
speaking, three cones: (1) the 
inner cone A, composed of com- 
bustible vapors; (2) an inter- 
mediate cone B, in which these 
vapors are decomposed by the 
heat and a 
small quantity 
of carbon is set 
free which ren- 
ders the flame 
luminous ; and (3) an almost invisible, nar- 
row outer cone, or film, C, in which the 
hydrogen and carbon are burned to water 
and carbon dioxide respectively. 

Bunsen burners. In the ordinary Bunsen 
burner used in chemical laboratories, and 
in similar burners used in gas stoves and 
ranges (Fig. 132) and for illumination with 
the aid of mantles, the gas is mixed with a 
certain percentage of air before it is burned. 
This is accomplished by having an open- 
ing (mixer) in the base of the burner 
(Fig. 132, A) into which the air is drawn 
by the flow of the gas. If the mixer is adjusted properly so 
that just the right amount of air is admitted, the flame is 
colorless. Such a flame possesses an advantage in that no 
carbon is deposited from it. 





Fig. 130. A match 
head suspended in 
lower part of gas 
flame is not ignited 






FUELS; ELECTKIC FUKNACES ; FLAMES 229 




Smoke prevention. Since the products of combustion of fuels 
are carbon dioxide and water vapor, and these are invisible 
compounds, it is evident that if the combustion is complete 
no smoke will be formed. As a rule 
the combustion is imperfect ; gaseous 
compounds containing carbon are first 
formed, and when these are imperfectly 
burned, a part of their carbon is set free in 
a finely divided state constituting s?noke. 
Smoke may therefore be prevented by 
securing the complete combustion of the 
fuel, the necessary conditions being as 
follows: (1) a sufficient supply of air; 
(2) thorough mixing of the air with the 
combustible gases produced from the 
fuel ; and (3) a temperature high enough 
a candle flame to maintain active combustion. 

Smoke prevention is a problem of great economic importance, 
especially in the large cities. Thus, for example, it has been 
estimated that the smoke in the 
city of Pittsburgh costs the people 
of the city $10,000,000 yearly, or 
about $20 for each inhabitant ; 
and this does not take into ac- 
count the serious effect of smoke 
upon health. Because of these 
facts many cities are now taking 
steps to abate the smoke nuisance. 
This is done by compelling the 
use of furnaces so constructed as 
to secure the conditions (noted 
above) for complete combustion. 

Gas mantles. In using the fuel gases as illuminants the gas is 
generally mixed with air before burning. In this way the gas 
burns with a hot but nearly nonluminous flame. The light is ob- 
tained by suspending about this flame a gauze mantle of suitable 




Fig. 132. A typical gas burner 



230 



CHEMISTRY AND ITS USES 



material. The successful development of the gas mantle was a great 
chemical achievement, for very few materials are suited to the pur- 
pose, and these are very rare in nature and difficult to purify. The 
best mantles are composed of a mixture of 99 per cent of thorium 
oxide with 1 per cent of cerium oxide. 

The thorium and cerium compounds used in gas mantles are 
obtained from monazite sand (Fig. 133), found principally in Brazil. 
The process of making a gas mantle consists in knitting a cylindri- 
cal cotton fabric, which is then dipped into a solution of the nitrates 




Fig. 133. Materials used in making gas mantles ; also different 
stages in the process 

of thorium and cerium. After drying, the fabric is heated, in which 
process the yarn is burned, while the nitrates of thorium and cerium 
are converted into oxides which are left in the form of the origi- 
nal fabric. The resulting mantle is very delicate and is strength- 
ened for shipping by dipping it into a solution of an appropriate 
substance and drying. 

EXERCISES 

1. Enumerate some important substances obtained in the destructive 
distillation of wood ; of coal. 

2. Why not use hard coal in preparing coal gas? 

3. How do you account for the presence of nitrogen («) in water gas 
(see table, p. 224) ? (b) in producer gas? 






FUELS; ELECTRIC FUKNACES ; FLAMES 231 

4. (a) Why do the windows of rooms heated by portable gas stoves 
sweat in cold weather? (b) Is the moisture on such windows on the 
inside or the outside of the glass ? 

5. Why do some samples of charcoal burn with a flame while others 
do not '? 

6. What rare element is found in the natural gas of certain localities ? 

7. What is the disadvantage of water gas as an illuminant? 

8. Which of the gases whose composition is given in the table 
(p. 224) should you expect to have the least heat value ? 

9. Venders of some forms of gas stoves often claim that no products 
are evolved when gas is burned in their stoves because the gas is com- 
pletely consumed. Is this contention correct ? 

10. Why is carbon deposited on cooking vessels heated over a flame ? 

11. A candle (composed of compounds of carbon and hydrogen) is 
burned in a closed volume of air. State the changes that take place 
both in' the candle and in the air. 

12. Why does bituminous coal burn with more of a flame than 
anthracite ? 

13. 1000 cu. ft. of methane weighs about 20.3 kg. Assuming natural 
gas to be pure methane, what is the heat evolved in the combustion of 
1000 cu. ft. of the gas ? 

14. The fuel value of a coal was determined by burning 1 g. of the 
coal in a calorimeter containing 2500 g. of water. The heat liberated 
raised the temperature of the water 3.1°. Calculate the calorific value 
of the coal. 

15. A portable gas stove used in heating a room burns 10 cu. ft. of 
gas each hour. If the gas is pure methane, (a) what volume of oxygen 
is withdrawn each hour from the air in the room ? (6) what volume of 
carbon dioxide is evolved? 



CHAPTER XXV 



COAL-TAR COMPOUNDS 



General discussion. We have seen that in the manufacture 
of coke and coal gas there is obtained a tarry mass known 
as coal tar. This is an exceedingly complex mixture of 
various compounds colored black with finely divided carbon. 




Fig. 134. View of a modern plant for distilling coal tar 

When coal gas was first prepared, this tarry mass was re- 
garded as a nuisance, and it was a source of expense to dis- 
pose of it. Later, chemists began to find in it some useful 
compounds, and this investigation has gone on until now 
coal tar is of inestimable value, furnishing us with thousands 
of compounds of the greatest importance to our welfare. 

232 



COAL-TAR COMPOUNDS 



233 



This achievement shows something of the part that the 
chemist is taking in the advancement of modern civilization. 
Usefnl products are obtained from coal tar by much the 
same methods as are used in refining petroleum. The tar is 
subjected to fractional distillation, and the various products 
which are separated in this way are then further purified. 




•CD gy ~J 



Im)Thracem£J 



Fig. 135. Some of the principal constituents obtained directly 
from coal tar 



tFig. 134 shows a modern plant for distilling coal tar. The large 
stills in which the coal tar is distilled are shown in the shed A, 
the vapors obtained on heating the tar are passed through con- 
densers (housed in B), and the resulting products are stored in 
containers, two of which, C, C, are shown in the figure. 
Coal-tar compounds. While there are many compounds in 
coal tar, it has been found economical to separate from it only 
eight or ten of them. The six most important of these are 
shown in Fig. 135. Each of these compounds, however, serves 
as the source material for the preparation of other compounds 
and these in turn for still others. All these compounds, whether 
present in coal tar or prepared from those that are present in 
coal tar, are known collectively as coal-tar compounds. Many 
thousands of them are known. 



234 CHEMISTKY AND ITS USES 

The chart (Fig. 136) gives us some idea of the kind of com- 
pounds that are derived from coal tar and also the relation of the 
compounds to each other. Thus, toluene is obtained directly from 
coal tar, as indicated by the arrow, while from toluene (following 
the arrows) we obtain benzoic acid, trinitrotoluene (T.N.T.), sac- 
charine, and Congo red. By simply following the arrows backward 
we can trace the ancestry of each compound named. Thus the dyes 
at the bottom of the chart are all derived from aniline, but aniline 
is prepared from nitrobenzene and this from benzene, which is 
obtained directly from coal tar. 

Properties and uses of important coal-tar compounds. It is 

possible to mention here only a few of the individual coal-tar 
compounds. We will simply note six compounds (Fig. 135) 
obtained directly from coal tar and a few typical compounds 
derived from each. 

1. Benzene (C Q H 6 ). This is a colorless, highly inflammable 
liquid boiling at 80°. When treated with nitric acid nitro- 
benzene (C 6 H 5 N0 2 ) is formed, and this, on reduction, yields 
a nearly colorless liquid known as aniline (C 6 H 5 NH 2 ), from 
which hundreds of dyes of all colors are prepared. 

2. Toluene (C 7 H S ). This resembles benzene in appearance. 
When treated with nitric acid it forms the solid compound 
trinitrotoluene (T.N.T.), so largely used as an explosive in the 
World War. From toluene there are also prepared saccharine, 
a compound five hundred times as sweet as sugar, and benzoic 
acid, whose sodium salt, sodium benzoate, is our most common 
food preservative. 

3. Carbolic acid (phenol) (CJSfiH). The pure acid is a 
white solid and is very poisonous. It serves for the prepara- 
tion of salicylic acid, from which we prepare the well-known 
medicine aspirin. Carbolic acid is also used in the prepara- 
tion of picric acid (a valuable explosive) and such substances 
as bakelite and condensite, used for making umbrella handles, 
pipestems, buttons, phonograph records, insulators in electrical 
devices, and many similar articles. 



SOME COAL PRODUCTS 

*« ,,£ — _ ► 



COKE. 






1 
1 T. N.T. 



COAL 

4r 



AMMONIA 



COAL 

(EXPLOSIVE) _ GAS 



BENZOIC ACID TOLUENE: 

(PRESERVATIVE), 



vasP 



TAR 



JIVE) p— I 

J Hi 



CARBOLIC ACID SAUCYUCAOD 



SACCHARINE 

(SWEEEMIN6 
A6EHT) 



u 

CONGO RED 

(DYE) 



r 



± 



BENZENE 



PICRIC ACID > f ASPIRIN 

(EXPLOSIVE) l_, (MEOiCiNAt) 






anthracene naphthalene 

u » 

ALIZARIN EOS1N 

(OYE) (OYE) 



fl CRESYUC AOD METHYL SALICYLATE 

N,TROBENZE ( f !NF ™ T) ( « *«> 



ANILINE HYDROQUINONE ACE1ANIIM 

(PHOTO DEVELOPER) (MEOrCINAL) 



I 



t 



BUTTER YaLOW $8 



1 



METHYL VmET (DYE) 



INDIGO 
(dye) 



(DYE) 



(DYE) 



ACIDVKH.ET 

(DYE) 



Fig. 136. Chart showing the derivation of some coal-tar compounds 
(Suggested by Science Service) 




Fig. 137. A case filled with American-made dyes 

The picture above the case is that of William Henry Perkin, the 
discoverer of aniline dyes 



COAL-TAR COMPOUNDS 235 

4. Cresylic acid {C^HftH). This substance is the basis of 
nearly all the common disinfectants used in our homes for 
such purposes as cleansing drainpipes and sinks. 

5. Naphthalene (C 1Q H S ). This is a white crystalline compound, 
best known to us in the form of white balls commonly called 
moth balls. 

6. Anthracene (C 14 2T 10 ). This is a nearly white solid and 
serves for the preparation of alizarin, one of our most valu- 
able dyes. 

Coal-tar dyes. All dyes derived from coal tar are known 
as coal-tar or aniline dyes. Many hundreds of these dyes have 
been prepared, of every imaginable color. It is interesting to 
note that two of the most common, namely, indigo and alizarin, 
were formerly obtained from vegetable sources. Chemists 
found out how to make them in the laboratory, and now they 
are made in this way at a much lower cost than formerly. The 
method of dyeing will be discussed in a later chapter. 

Previous to the World War nearly all our dyes came from 
Germany. During the war this supply was cut off, and there 
was a temporary shortage of dyes. American chemists soon 
solved the problem connected with their production, and now 
the United States produces an adequate supply of dyes of the 
very highest grade (Fig. 137). 

Historical. The first aniline dye was made in 1856 by an Eng- 
lish boy seventeen years of age. His name was William Perkin. 
The boy was assisting in the laboratory of an English university, 
and during the holidays he spent his time trying to make quinine. 
In the course of some experiments with aniline he noticed that a 
colored substance of great beauty was produced. At this time all 
the dyes used were obtained from vegetable sources. Perkin got 
the idea that perhaps this new compound which he had prepared 
from aniline might be used as a dye. He finally succeeded in 
showing that it could be used and that it was superior to vegetable 
dyes in coloring power. This discovery led to others, and the 
investigation still continues, as shown by the fact that our dye 



236 



CHEMISTKY AND ITS USES 



plants employ hundreds of chemists. It is interesting to note that 
in 1906, fifty years after Perkin's discovery, he came to the 
United States and attended a great meeting held in New York 
City in honor of the fiftieth anniversary of the discovery of the 
first aniline dye. 

Coal-tar compounds in foods. Much discussion has arisen 
over the use of coal-tar compounds in foods. The Federal 
government 



has selected eight aniline 




dyes of different 
colors the use of 
which is permitted 
in such foods as can- 
dies and butter. As 
already stated, the 
use of sodium ben- 
zoate as a preserva- 
tive is allowed under 
certain restrictions. 
Saccharine was for- 
merly permitted in 
foods, but in 1912 
the government for- 
bade the further use 
of it. Vanillin, identical with the compound prepared from 
vanilla beans, and coumarin, which has an odor similar to vanil- 
lin, are both used in artificial vanilla extracts, but when they 
are so used the label on the container must state the fact. It 
is well to keep in mind that all such substances are not foods 
and are used for purposes other than nutrition. 

Coal-tar compounds in medicines. Coal tar also furnishes us 
with many useful medicines. Among these are acetanilide, 
prepared by treating aniline with acetic acid and used in head- 
ache powders ; procaine, a valuable local anesthetic used in 
minor surgical operations and in extracting teeth ; stovaine, also 
a local anaesthetic ; salicylic acid and its derivative aspirin. 



Fig. 138. Using pitch obtained from coal tar in 
road construction 



COAL-TAR COMPOUNDS 237 

One of the most valuable of all these medicines is the com- 
pound known as arsphenamine (salvarsan). Many of our 
chemists are now working in the development of new com- 
pounds which promise to be of great service in the treatment 
of certain diseases that are now incurable. 

Pitch and its uses. When coal tar is distilled there remains 
in the still a very thick, viscous, tarry mass known as pitch, 
which has extensive uses in road construction. Tar is fre- 
quently sprinkled on roadways to serve as a binder and thus 
prevent the wearing away of the road (Fig. 138). 

EXERCISES 

1. Are all the so-called coal-tar compounds present in coal tar? 

2. («) What is the difference in composition between benzene and 
benzine ? (7>) Have the two substances any properties in common ? 

3. Do you see any reason why benzene could not be used in place of 
gasoline in a gasoline engine ? 

4. Trace the steps in obtaining aspirin from coal tar (Fig. 136). 

5. Mention four important compounds prepared from toluene and 
the uses of each. 

6. Trace the steps in obtaining indigo from coal tar. 

7. Mention five different compounds obtained from coal tar, each 
used for a different purpose, and state the use of each (Fig. 136). 

8. Mention a dye and an explosive, both of which are prepared from 
the same compound. Does this suggest a reason why the Germans used 
their dye factories for making explosives during the World War ? 

9. What do you think of the propriety of allowing foods to be colored 
artificially ? 

10. Calculate what weight of benzene will be required to make 
100 kg. of aniline, assuming that 1 gram-molecule of benzene will give 
1 gram-molecule of aniline, thus : 

>C 6 H 5 NH 2 



CHAPTER XXVI 

CARBOHYDRATES AND TEXTILES 

Carbohydrates. The term carbohydrate is applied to a class 
of compounds which includes the sugars, starch, and allied 
substances. These compounds contain carbon, hydrogen, and 
oxygen, the last two elements usually being present in the 
proportion in which they combine to form water. The most 
important carbohydrates are the following: 

TABLE OF CARBOHYDRATES 

Sucrose (ordinary sugar) C 12 H 22 O n 

Lactose (milk sugar) C 12 H 22 O n 

Maltose C 12 H 22 O n 

Dextrose (grape sugar) C 6 H 12 6 

Levulose „ C 6 H 12 6 

Dextrin (C 6 H 10 O 5 ) x 

Cellulose (C 6 H 10 O 5 ) x 

Starch (C 6 H 10 O 5 ) x 

The molecular formulas of dextrin, cellulose, and starch are 
unknown, but are multiples of the simple formula C 6 H 10 O 5 ; 
hence they are often written (C 6 H 10 O 6 ) x . In the discussion of 
the compounds they will be represented by the formula C 6 H 10 O 5 . 
It will be noted that some of the compounds named in the 
above table have the same formula. Compounds having the 
same formula are said to be isomeric. The difference in the 
properties of such compounds is due to the fact that the atoms 
are arranged differently in the molecule. Just as two houses 
may be entirely different and yet each be built from the 

238 



CARBOHYDRATES AND TEXTILES 



239 



same kind and the same number of brick, so two compounds 
may be different, and yet the molecule of each contain the 
same elements and the same number of atoms of each. 

Sucrose (sugar) (C 12 H 22 O u ). This substance, commonly 
called sugar, occurs in many plants. The world's supply comes 
entirely from the sugar cane and the sugar beet, the sugar cane 




Fig. 139. Source of the world's output of sucrose, 
and the sugar beet 



the sugar cane 



furnishing about 70 per cent of the total output. The sugar 
cane grows only in warm climates (Cuba and the Hawaiian 
Islands are the greatest producers) ; the sugar beet thrives in 
cooler climates (Fig. 139), such as prevail in Germany or in 
Ohio and Michigan in the United States. The beets contain 
as high as 16 per cent of sucrose. 

The separation and refining of sugar. To obtain the sugar from 
the sugar beet, the beets are sliced and then placed in large cylin- 
drical containers. Water is run through these containers and 
dissolves the sugar in the beets. Unfortunately it also dissolves 



240 CHEMISTRY AND ITS USES 

many other substances present, so that the chemist must find ways 
for removing these impurities. This is done partly by precipitation 
and partly by filtering through some porous material such as bone- 
black. The resulting solution is then evaporated in closed vessels 
(vacuum pans A, Fig. 140) from which the air is partially exhausted. 
In this way the boiling point of the solution is lowered and the 
charring of the sugar prevented. After the sugar crystallizes out 
it is separated from the remaining sirup by transferring it to 
perforated vessels which are whirled at a high rate of speed 
(centrifugals) (Fig. 141). The sugar is separated from the juice 
of the sugar cane by making use of the same general principles. 
The impurities in the beet juice are more distasteful than those 
in the cane juice, hence all beet sugar must be highly refined. 
Ordinary brown sugar is cane sugar only partially refined. It is 
not practicable to separate all the sugar from its solution. The 
sirup that remains when sugar is prepared from sugar cane con- 
stitutes ordinary molasses. The sweetness of maple sugar and 
maple sirup is due to sucrose, other products present in the maple 
sap imparting the distinctive flavor. Cane sugar and beet sugar, 
when pure, are identical, and even the chemist cannot distinguish 
between them. The annual consumption of sugar in the United 
States amounts to nearly 100 lb. for each person. 

Chemical conduct of sucrose. When a solution of sucrose is 
heated to about 70° with water and hydrochloric acid, the 
sugar and water react to form two isomeric sugars, dextrose 
and levulose, as shown in the equation 

W- + H »°^ C M + c 6 h 12 o 6 

In this process the sugar is said to be inverted, and the mixture 
of dextrose and levulose is called invert sugar. These two 
sugars are white solids. They are not so sweet as sucrose, 
neither do they crystallize so easily ; hence in making candy 
from sucrose it is customary to add a small amount of acid 
(vinegar) or acid salt (such as cream of tartar), which produces 
some invert sugar, whose presence prevents the remaining 
sucrose from crystallizing. 





Fig. 140. Vacuum pans in a sugar factory 
The air is pumped from the pans A hy the engine B 




Fig. 141. Centrifugals for removing the sugar from sirup 




Fig. 142. Cheese stored for " ripening " 




Fig. 143. Making glucose (corn sirup) 

Starch and a little hydrochloric acid are heated with steam in a large 
copper vessel A, which is known as a " converter " 






CARBOHYDRATES AND TEXTILES 241 

When heated to 160° sucrose melts ; if the temperature is in- 
creased to about 215° a partial decomposition takes place and 
a brown substance known as caramel is formed. This is used 
extensively as a coloring matter and in making confectionery. 

Lactose (milk sugar) (C 12 H 22 11 ). This compound is a white 
solid, but is not so sweet or so soluble as sucrose. It is present 
in the milk of all mammals. The average composition of cow's 
milk is as follows : 

Water 87.17% 

Casein (nitrogenous matter) 3.56% 

Butter fat 3.64% 

Lactose 4.88% 

Mineral matter 0.75% 

The casein is a solid and is present in sweet milk in a finely 
divided state. When the milk sours, the casein separates. 
The remaining liquid, known as whey, contains the lactose, 
which can be obtained by evaporation. The souring of milk 
is due to the fact that the lactose present changes into lactic 
acid, a liquid having the formula C 3 H 6 3 . 

C 12 H 2 Ai + H 2 0— *-4C 3 H 6 3 

This change is brought about through the agency of a certain 
microorganism which enters from the air, and the process is 
known as lactic fermentation. The body of the ordinary medi- 
cine tablet consists of lactose because this substance readily 
absorbs medicinal solutions. 

The souring of milk. It has been stated that the souring of 
milk is caused by a certain microorganism. This organism is 
really a microscopic plant which grows in the milk, and in so 
doing changes the lactose into lactic acid. The plant does not 
grow at low temperatures ; hence the souring of milk is pre- 
vented (or delayed) by keeping it cold. Extreme cleanliness 
also tends to retard souring by preventing the introduction 
of the plant, the spores of which are present in filthy matter. 



242 



CHEMISTRY AND ITS USES 



The pasteurization of milk. Certain diseases may be spread by 
impure milk just as by impure water (p. 66). The microorganisms 
which cause these diseases may be killed by boiling the milk as 
in the case of water. This process, however, causes changes in the 
milk which impair its food value. The same results can be accom- 
plished, without impairing the food value of the milk, by heating 
it to 55°-70° for several hours. Such milk is said to he, pasteurized, 

the term being 
derived from the 
name of the great 
French chemist 
Pasteur, who first 
showed that many 
changes and dis- 
eases are due to 
microorganisms. 

The making of 
cheese. Cheese 
is made from 
the casein of 
milk. Ordina- 
rily the fat of the milk is left in with the casein (cream cheese). 
To separate the casein there is added to the milk a substance 
known as rennet, which is extracted from calves' stomachs. 
When the rennet is added to the milk it causes the casein 
to separate (or set) in the form of a porous solid, which con- 
tains the other ingredients of the milk in its pores. After the 
casein sets (Fig. 144) it is cut by knives and separated from 
the whey. The resulting casein is then salted, compressed in 
molds, and set away in a cool place for several weeks to ripen 
(Fig. 142). In this process certain changes take place which 
produce small percentages of compounds which impart to the 
cheese its characteristic taste. These changes are brought 
about by microorganisms, and each kind of organism produces 
definite changes, By inoculating the newly made cheese with 




Fig. 144. When rennet is added to milk the casein 
separates (or sets) in the form of a porous solid 



CAEBOHYDEATES AXD TEXTILES 



243 



the proper microorganism it is possible to secure the desired 
flavor, or, in other words, to prepare the kind of cheese desired. 

Maltose. When certain grains, like barley, start to grow, the 
sprouting grains develop a substance known as diastase, which 
has the power of changing starch into a sugar known as maltose. 
Malt (Fig. 145) consists of barley grains which have been kept in 
a moist, warm place until the grains germi- 
nate and have then been heated to destroy 
the vitality of the germ. The malt contains 
diastase and is used for converting starch 
into maltose in the preparation of alcohol. 
The diastase acts as a catalytic agent. 

Dextrin. This is a sweetish solid obtained 
by roasting starch under proper conditions ; 
hence it is present in toast and bread crust. 
It is soluble in water and has many uses, 
such as making gum, mucilage, and paste. 
It is used on the back of postage stamps and 
stickers of all kinds. 

Dextrose (grape sugar) (CLKoO V This 

. • -, 6 i • Fig. 145. Malt used 

sugar is present in honey and m many . . .. 

o i j j m preparing maltose 

fruits, usually associated with levulose, 
and is often called grape sugar, because of its presence in 
grape juice. It can be obtained, along with levulose, by heating 
sucrose with hydrochloric acid, as explained on page 240. 
Commercially it is prepared by heating starch and water with 
a small percentage of hydrochloric acid. 




C e H O +H,0 



CJI O 



When the change is complete the hydrochloric acid is neu- 
tralized by sodium carbonate and the solution is evaporated. 

Glucose (corn sirup). When starch and water are heated with 
hydrochloric acid (catalytic agent) pure dextrose is obtained, as 
stated above. If the process is stopped short of completion the 
product consists of dextrose, maltose, and dextrin, sufficient water 



244 



CHEMISTRY AND ITS USES 



being present to form a thick, colorless sirup. This product is 
known as glucose. Enormous quantities of it are prepared. The 
starch used in its preparation is obtained from corn, about 65,000,000 




Fig. 146. Removing the starch from the settling troughs in 
a starch factory 

bushels of corn being used yearly for this purpose. The starch, 
water, and acid are introduced into large copper retorts (A, Fig. 143) 
fitted with steam pipes and heated under pressure with live steam. 

Glucose is much cheaper than sugar 
and is used chiefly in making candy, 
jellies, jams, and table sirups. One 
brand of corn sirup, largely used, con- 
sists of 85 per cent of glucose and 
15 per cent of molasses, the latter 
being added to impart the flavor. 




Fig. 147. Wheat-starch gran- 
ules magnified 200 diameters 



Starch (C 6 H 10 O 5 ). This substance 
is always present in seeds and 
tubers and is one of the abundant 
carbohydrates found in nature. In 
the United States it is obtained chiefly from corn, about 60 per 
cent of which is starch; a limited supply is obtained from 
potatoes. In Europe, on the other hand, the potato serves 
as the principal source. 



CARBOHYDRATES AND TEXTILES 



245 




Manufacture of starch. In manufacturing starch, from corn, the 

corn is first soaked in water containing a little sulfurous acid to 

soften the grain. It is then ground coarsely so as not to crush 
the germ. When the resulting mass is 
mixed with water the germ floats, being 
very light because of the oil which it con- 
tains. In this way the germ is separated 
from the rest of the seed, and from it corn 
oil is prepared. The remaining material, 
consisting of the starch, the nitrogenous 
constituent (gluten), and the bran, or out- 
side coating of the grain, is then ground 
fine, mixed with water, and passed through 
cloth sieves, which remove the bran. The 
water containing the starch and gluten in 

suspension is then allowed to run slowly down long, shallow 

troughs, the rate of flow being regulated so that the , heavier 

starch sinks to the bottom of the 

trough while the lighter gluten is 

washed away. The starch is then 

removed from the troughs (Fig. 146) 

and dried. Large quantities of 

starch are used in making glucose 

and other foods, for finishing cloth, 

and for laundry purposes. 



Fig. 148. Granules of 

cornstarch magnified 

200 diameters 




Characteristics of starch. Starch 
consists of minute granules which 
differ somewhat in appearance, 
according to the source of the 
starch, so that it is often possible 
from a microscopic examination to 
determine from what plant any 

given sample of starch was obtained (Figs. 147 and 148). 
When heated with water the granules burst and the starch 
partially dissolves. This is the reason why starchy foods are 
made more digestible by cooking. 



Fig. 149. Two very important 
derivatives of cellulose 



246 



CHEMISTRY AND ITS USES 




mubbhw mm 

R51 



Cellulose (C 6 H 10 5 ). Cellulose is the basis of all wood fibers. 
Cotton and linen are nearly pure cellulose. Cellulose is in- 
soluble in water and dilute acids, but dissolves in a solution 
prepared by dissolving copper oxide in am- 
monium hydroxide. Concentrated hydro- 
chloric acid changes it into dextrose. A 
mixture of sulfuric and nitric acids acting 
upon it forms a number of compounds col- 
lectively known as nitrocellulose or guncotton 
(Fig. 149). These are white solids, but they 
can also be obtained in a transparent form 
(Fig. 150). They are inflammable and, under 
certain conditions, highly explosive. They 
are the chief constituents of most smokeless 
powders. Photographic films (Fig. 150) are 
made from them, as well as from a derivative 
of cellulose known as cellulose acetate, which 
is not so inflammable as nitrocellulose. 
Collodion is a solution of nitrocellulose in 
a mixture of alcohol and ether, Celluloid is 
a mixture of nitrocellulose and camphor (a 
white solid obtained from the camphor tree, 
which grows in Japan). These two, when 
mixed together, form a plastic mass which 
can be molded into any desired shape and 
which is used for making such objects as 
combs and brush handles. When such arti- 
cles are warmed by holding them between 
the hands, the odor of camphor is easily 
detected. Nitrocellulose is also used in mak- 
ing a sort of artificial leather (fabrikoid),- 
from which handbags and trunks are made. 

Natural silk. Silk has long been known and highly prized, as is 
shown by the references to it in the Bible. The silk fiber is spun by 




Fig. 150. Section of 

movie film made of 

nitrocellulose 



CARBOHYDRATES A]S T D TEXTILES 



247 



the silkworm. These are grown in large numbers (Fig. 151), espe- 
cially in Japan and China, and utilized in the production of silk. 
Mercerized cotton and artificial silk. When cotton cloth is treated 
with a concentrated solution of sodium hydroxide, the cellulose 
shrinks and becomes tougher in character. If the cloth is placed 




Fig. 151. The production of natural silk 

The silk is spun by the silkworm. The worm (B, B, B) feeds best on mulberry 
leaves. When ready to spin it fastens itself by threads to the branches of the 
tree (Q and spins silk fibers (often 1000 yards in length) about its body, forming 
the cocoon (D, E). In this process the worm is gradually transformed into the 
chrysalis (F), then into the insect, which breaks through the cocoon (G) and 
becomes free. .The insect (^1) lays the eggs from which the worm is hatched, 
and the cycle repeats itself. The threads forming the cocoons are wound onto 
reels and are used for weaving silk fabrics 



248 



CHEMISTRY AND ITS USES 



in stretchers to prevent the shrinkage, it assumes an appearance 
somewhat resembling silk and is known as mercerized cotton. 
Another fabric prepared in large quantities from cellulose resem- 
bles silk very closely and is known as artificial silk (Fig. 152). 




Fig. 152. The production of artificial silk (viscose) in a French factory 

Cellulose or one of its derivatives is dissolved, and the concentrated solution is 
forced through minute tuhes and received into appropriate liquids which cause 
the cellulose to coagulate in the form of fine threads. France is a leader in the 
production of artificial silk, hut the United States also produces a large amount. 
Artificial silk resembles natural silk in appearance, but is not so strong. (From 
Duncan's " Chemistry of Commerce ") 

Characteristics of various textile fibers. Of the different 
fibers used in making the yarns from which the common 
fabrics are prepared, the vegetable fibers, cotton and linen, 
are essentially cellulose, while the animal fibers, wool and silk, 
are composed of nitrogenous substances. Although these fibers 
resemble each other when viewed with the naked eye, their 






CARBOHYDRATES A^T> TEXTILES 



249 






appearance varies widely when examined with the microscope 
(Fig. 153). It is also possible to distinguish between the fibers 
by the action of chemical reagents. For example, a hot solu- 
tion of sodium hydroxide (5 per cent to 10 per cent) has but 
little action upon cotton, while it will readily dissolve wool 
and slowly dissolve silk. The nitrogenous fibers burn with 
a foul odor like that of burning hair, and this property is 
sometimes useful in determining the source of a fiber. 






Silk fiber 



Cotton fiber 



Wool fiber 



Fig. 153. Textile fibers (highly magnified) 

Paper. Paper consists mainly of cellulose, the finer grades 
being made from linen and cotton rags and the cheaper grades 
from wood. 

Manufacture of paper. Most of our paper is made from wood. 
The trees, such as spruce and hemlock, are cut down, floated down 
to the mills, and there" sawed into short lengths. The bark is 
then removed and the wood cut into small chips by allowing the 
short lengths to slide down a trough against a rapidly revolving 
wheel fitted with sharp blades. The resulting chips are then 
placed in large retorts and heated either with a solution of sodium 
hydroxide or with calcium acid sulfite (a compound prepared by 
passing sulfur dioxide upwards through large towers filled with 
limestone, over which water trickles); these dissolve the objection- 
able matter from the wood, leaving the cellulose, which is then 
bleached with chlorine and washed with water. The cellulose is 
thus obtained in the form of thin shreds suspended in water, and 
this is known as paper pulp. This pulp (A, Fig. 154) is run onto 
wire screens, and most of the water present runs through the 
screen. It then passes between large iron cylinders, some of 



250 CHEMISTRY AND ITS USES 

which are heated with steam. In this way the pulp is pressed and 
dried and delivered in the form of paper (B, Fig. 154). In the 
process different materials are often added to the pulp. These 
vary with the nature of the paper desired; thus finely ground 
clay or calcium sulfate is added to give body to the paper. In 
making paper intended for writing or printing, a compound pre- 
pared by heating resin and sodium hydroxide is added, together 
with aluminium sulfate. This prevents the ink from spreading. 







~— — — -__ 




!| ^— — - 




■■■-■—- — ... 


~~~~— ^ 




^-^Jr 3 : 


>— rl «^r»»w 


S ==CT^ 


mmjjjgpzm 












■ : " ; .- 




B 




™§^1^| 


i&mzm 


"■ " : .""' ■ 












^mg^g 




\j->_ 
















\. 









Fig. 154. View of the interior of a paper mill 

The alkaloids. We have seen in this chapter that plants 
furnish us with food and clothing. They also furnish us with 
many useful medicines. Prominent among these is the class 
of substances known as alkaloids because of their basic proper- 
ties. Each of these is found in some part of a certain plant. 
Thus caffeine is always present in the coffee berry and tea 
leaves and is their active constituent. Some of the most 
important alkaloids are quinine, morphine, strychnine, and 
cocaine. All the above are white, solid compounds containing 
carbon, hydrogen, oxygen, and nitrogen. Nicotine, the active 
constituent of tobacco, is one of the few liquid alkaloids. 

The alkaloids, as a rule, are very active, and most of them 
are intensely poisonous. Some of them, especially cocaine 
and morphine, are very dangerous habit-forming drugs. 






CARBOHYDRATES AND TEXTILES 251 

EXERCISES 

1. What is the relation between the number of atoms of hydrogen 
and oxygen in each of the carbohydrates given in the table on page 238. 

2. Distinguish between the two terms allotropic and isomeric. 

3. (a) Is there any difference between unrefined cane sugar and unre- 
fined beet sugar ? (b) between refined cane sugar and refined beet sugar ? 

4. It is sometimes stated that milk sours more readily during 
thunderstorms. Do you think there is any evidence for the truth of 
this statement? 

5. What constituents of milk are used in the making of cheese? 

6. To what is («) the sweet taste of milk due ? (6) the sour taste ? 

7. Why do different kinds of cheese have different flavors ? 

8. What is the function of the hydrochloric acid used in making- 
glucose ? 

9. Is any constituent of toasted bread soluble in water? 

10. Enumerate the uses of cellulose. 

11. How could you distinguish (a) between artificial and natural 
silk? (b) between mercerized cotton and silk? 

12. Celluloid is readily inflammable. What is the combustible con- 
stituent present in it ? 

13. Can you suggest any chemical change that may take place when 
starched clothes are ironed? 

14. What weight of starch is necessary for the preparation of 100 kg. 
of dextrose ? 

15. What weight of invert sugar can be prepared from 100 g. of 
sucrose ? 



CHAPTER XXVII 



ALCOHOLS ; PRESERVATIVES 

The alcohols. Just as there are a great many sugars, so 
there are a great many alcohols. The two most important 
ones are methyl alcohol (CH 3 OH) and ethyl alcohol (C 2 H g OH). 
The latter compound is the common alcohol. Both of these 
alcohols are colorless, inflammable liquids. 

Methyl alcohol and ethyl alcohol may he regarded as derived 
from methane (CH ) and ethane (C 2 H 6 ) by substituting an OH? 
group for an atom of hydrogen ; hence the terms methyl and ethyl. 

Methyl alcohol (wood alcohol) (CH 3 OH). This compound is 
formed when wood is heated in the absence of air (p. 106), 
and on this account it is called wood alcohol. It is a colorless 
liquid which boils at about 66° and burns with an almost 
colorless flame. It is a good solvent for organic substances 
and is used extensively in the manufacture of varnishes. It 
is quite poisonous, and many deaths occur from its use as a 
substitute for ethyl alcohol. It acts upon the optic nerve, and 
many persons have become blind from drinking the liquid or 
from repeatedly inhaling its vapor. 

Duncan states that " out of 10 men who drink 4 oz. of pure 
methyl alcohol in any form whatever, 4 will probably die, 2 of 
them becoming blind before death. The remaining 6 may recover, 
but of these, 2 will probably be permanently blind." 

When a mixture of the vapor of methyl alcohol and air is 
passed over hot copper, the alcohol is partially oxidized, forming 
a gaseous compound known as formaldehyde : 

2 CH 3 OH + 2 y 2 CH 2 + 2 H 2 

252 



ALCOHOLS ; PKESEKVATIVES 



253 



This gas is now prepared in large quantities and used as a disin- 
fectant and also to prevent decay. A 40 per cent aqueous solution 
of it is sold under the name formalin. 

Ethyl alcohol (grain alcohol, alcohol) (C 2 H 5 0H). This com- 
pound is the one commonly known as alcohol. It is a colorless 
liquid with a pleasant odor and is an excellent solvent for 
many organic substances. It boils at 78.3°. It is sometimes 
used as a fuel, since its flame is very hot and does not deposit 
carbon, as the flame from oil does. When taken into the system 
in small quantities it causes intoxication ; in larger quantities 
it acts as a poison. The intoxicating properties of such liquors 
as wine and whisky are due to the alcohol present. The ordi- 
nary alcohol of the druggist contains about 95 per cent of 
alcohol and 5 per cent of water. A solution containing 99 per 
cent or more of alcohol is called absolute alcohol. When alco- 
hol is heated with sulfuric acid a low-boiling inflammable 
liquid known as ether is formed : 



2 C 2 H 5 OH 



This is largely used as an anaesthetic 
in surgical operations. 

Preparation of ethyl alcohol. Al- 
cohol is prepared by the action of 
ordinary baker's yeast upon differ- 
ent sugars such as dextrose : 




C,H 12 6 



2 C 2 H 6 OH + 2 C0 2 



Fig. 155. Some cells of the 
yeast plant (magnified) 
This process in which a sugar is 

changed into alcohol and carbon dioxide by the action of 

yeast is known as alcoholic fermentation. The yeast is a low 

form of plant life (Fig. 155) and thrives in appropriate sugar 

solutions. During its growth a number of changes take place 

which result in converting the sugar into alcohol and carbon 

dioxide, as is shown in the above equation. 



254 CHEMISTRY AND ITS USES 

Alcohol can also be prepared by adding yeast to a solution 
of sucrose. The sucrose is first changed into dextrose and 
levulose (invert sugar, p. 240) by the yeast. The dextrose 
and levulose then ferment. 

Commercial preparation of alcohol. Most of the alcohol made in 
the United States for commercial purposes is prepared from crude 
molasses, which is brought from Cuba in large quantities for this 
purpose and directly fermented. A limited amount is made from 
starch obtained from corn. When starch is used it is first mixed 
with malt (p. 243), which changes the starch into maltose. Yeast 
is then added, and the maltose ferments. The alcohol is separated 
by fractional distillation. 

The alcohol problem. From very ancient times men of all 
nations have prepared alcoholic liquors of various kinds as bev- 
erages. In alcoholic content these range from 3 to 5 per cent 
in beers and light wines to 50 per cent and higher in whisky 
and brandy. Governments have found these beverages to 
be a fruitful source of revenue and have imposed heavy 
taxes on all kinds of fermented and distilled liquors, care- 
fully supervising their manufacture. They have also taxed 
pure alcohol heavily because imitation beverages can be made 
very easily from pure alcohol. In this country the Federal 
government imposed a tax on 95 per cent alcohol, which 
before the war was $2.10 per gallon and since the war $4.20 
per gallon. 

On the other hand, alcohol has important scientific and indus- 
trial uses for which no good substitute can be found, and the 
government has no desire to interfere with these uses, which 
the Federal prohibition act defines very carefully. Accordingly, 
educational and scientific institutions can obtain alcohol tax 
free for scientific uses, and for industrial purposes the govern- 
ment allows the use of denatured alcohol free of tax. Denatured 
alcohol is alcohol to which has been added a definite amount 
of some substance which renders it unfit for use as a beverage 



ALCOHOLS ; PRESERVATIVES 



255 



but which does not impair its use for manufacturing pur- 
poses. The substances which may be added for this purpose 
are prescribed by law and are known as denaturants. The 
most common denaturant is pyridine, a vile-smelling # sub- 
stance obtained by heating bones in the absence of air. 




Fig. 156. Interior view of an industrial alcohol plant 
Distilling off the alcohol 

Soft drinks (pop). The soft drinks on the market consist of 
water charged with carbon dioxide, sweetened with sugar or glu- 
cose, and flavored with various substances such as vanillin and 
the oils obtained from lemon and orange peeling. Oftentimes they 
contain a small amount of citric or tartaric acid (pp. 261, 262). 
A few of them contain caffeine (p. 250), which is the active 
constituent of tea and coffee. The colored ones contain aniline 
dyes, as can easily be proved by heating a highly colored pop 
for five or ten minutes with a strip of white woolen fabric. The 
dye leaves the liquid and dyes the wool a brilliant color. 



256 CHEMISTRY AND ITS USES 

Enzymes. We will recall that malt (p. 243) contains a 
substance known as diastase, which has the property of chang- 
ing starch into a sugar. This diastase belongs to a class of 
substances known as enzymes. There are many enzymes, and 
they often are of great importance. Thus yeast produces the 
enzyme known as zymase, and it is the zymase which really 
changes the sugar into alcohol and carbon dioxide, ifeast, 
therefore, is a sort of manufacturing plant for the production 
of zymase. It also produces an enzyme known as invertin, which 
changes sucrose into invert sugar. Enzymes play an important 
part in the process of digestion. For example, the active prin- 
ciple of the gastric juice is the enzyme pepsin. We know little 
about the chemical nature of enzymes. They are produced by 
growing organisms and seem to act as catalytic agents. 

Chemical changes in bread-making. The average composi- 
tion of wheat flour is as follows: 

Water 11.9% 

Gluten (nitrogenous matter) 13.3% 

Fats 1.5% 

Starch 72.7% 

Mineral matter 0.6% 

In making bread, flour is mixed with water, yeast, and a 
little sugar, and the resulting dough is set aside in a warm 
place for a few hours. The yeast first causes the sugar to 
undergo alcoholic fermentation. The carbon dioxide formed 
escapes through the dough, making it light and porous. The 
yeast plant thrives best at about 30° ; hence the necessity for 
keeping the dough in a warm place. In baking bread the heat 
expels the alcohol and also expands the bubbles of carbon 
dioxide caught in the dough, causing it to become porous 
and making the bread light. 

Preservatives. We have observed that the changes taking 
place in the souring of milk and the changing of sugar into 
alcohol are caused by microorganisms, the cells of which are 



ALCOHOLS ; PRESERVATIVES 



257 



present in the air. Many other similar changes, such as putre- 
faction (decay), are due to the same causes. All these changes 
may be prevented in one of the following ways: 

1. By keeping the substance at such a low temperature that 
the organism causing the change cannot thrive (cold storage). 

2. The substance may be heated so as to destroy all organ- 
isms present and then sealed air-tight in a suitable container. 
This is the method used in canning vegetables. Foods treated 
in this way are said to 
be sterilized. 

3. Some substance 
may be added which 
in small amounts will 
destroy the organisms 
causing the change 
or will prevent their 
growth. Such a sub- 
stance is known as a 
preservative. 

Whether or not pre- 
servatives should be 
permitted in foods is a much debated question. Some people 
maintain that any substance which is powerful enough to 
prevent the growth of the organisms must have an injurious 
action upon digestion. The Federal government at present 
allows the use of sodium benzoate (p. 234) in such foods as 
jellies, jams, and catchup, which are not consumed immedi- 
ately upon the opening of the container. If this preservative 
is used, however, the labels on the containers must state the 

amount present. 

• 

Fig. 157 shows the method used in sterilizing canned corn in 
a large factory. The cans are filled with corn and are then placed 
in a large wire basket A, of such a size as to fit into the retorts 
(1, 2), a number of which are shown in the figure. When the 




Fig. 157. Sterilizing- canned corn 



258 CHEMISTRY AND ITS USES 

retort is filled the lid is clamped on and the corn heated by lead- 
ing live steam into the retort. A small opening is left in the cans 
during the heating so as to allow for the escape of air. When the 
corn has been heated long enough to destroy all microorganisms, the 
cans are removed and the opening sealed under conditions that 
prevent the admission of any air. 

EXERCISES 

1. How could you distinguish between methyl alcohol and ethyl 
alcohol ? 

2. What is the name of the process used in separating alcohol 
from water? 

3. Give the changes involved in making alcohol from corn. 

4. Why is a little sugar or molasses added in making bread? 

5. For what purpose is yeast always used? 

6. Distinguish between lactic fermentation and alcoholic fer- 
mentation. 

7. Why does the government permit the use of a preservative 
(sodium benzoate) in tomato catchup but not in milk? 

8. Suggest a method for making alcohol from sawdust. 

9. What compounds have we studied that are used as anaesthetics? 

10. It has been proved that automobiles may be run by alcohol in 
place of gasoline. What is the objection to its use for this purpose ? 

11. (a) What are the products of combustion of alcohol? (b) Write 
the equation for the reaction that takes place when it burns. 

12. (a) Is alcohol formed in making bread? (b) Does the finished 
product contain any alcohol ? 

13. How could you prevent (a) sweet cider from becoming hard? 
(b) grape juice from becoming wine ? 

14. What weight of dextrose is necessary for the preparation of 
100 kg. of the alcohol of the druggist, assuming that 95 per cent of the 
dextrose ferments ? 

15. What weight of pure methyl alcohol is necessary for the prepara- 
tion of 1 kg. of formalin ? (Formalin contains 40 per cent by weight 
of formaldehyde.) 



CHAPTER XXVIII 

ORGANIC ACIDS AND THEIR DERIVATIVES; PROTEINS 

Organic acids. A great number of acids are known which 
are composed of carbon, oxygen, and hydrogen, and as a group 
these are called organic acids. Like the hydrocarbons, they 
can be arranged in series, one of the most important of which 
is known as the fatty-acid series. A few of the most important 
acids of this series are given in the following table. They are 
all monobasic — a fact indicated in the formula by separating 
the replaceable hydrogen atom from the rest of the molecule. 



SOME FATTY ACIDS 

CH0 2 formic acid, a liquid boiling at 100 c 

C 2 H 3 2 acetic acid, a liquid boiling at 118° 

C 4 H 7 2 butyric acid, a liquid boiling at 163 c 

C 16 H 31 2 palmitic acid, a solid melting at 62 c 

C 18 H 35 2 stearic acid, a solid melting at 69° 

C n H 2n _ 1 2 general formula 



Of these acetic acid deserves special mention. 

Acetic acid (H • C 2 H 3 2 ). This is the acid which imparts 
the sour taste to vinegar. It is prepared commercially by the 
destructive distillation of wood (p. 106). It is a colorless liquid 
and has a strong, pungent odor. When anhydrous it crystal- 
lizes as a white solid which melts at 18° and closely resembles 
ice in appearance ; hence the name glacial acetic acid. Many 
of the salts of acetic acid (acetates) are well-known compounds. 
Thus lead acetate (Pb(C 2 H 3 2 ) 2 - 3 H 2 0) is the white solid 
known as sugar of lead. 

259 



260 



CHEMISTRY AND ITS USES 



Vinegar. As is well known, when cider is exposed to the air it 
is gradually transformed into vinegar. Two changes are involved 
in the process : (1) the sugar in the cider first undergoes alcoholic 
fermentation, forming hard cider, which contains from 4 to 8 per 
cent of alcohol ; (2) the alcohol is then oxidized to acetic acid, the 
necessary oxygen coming from the air. This oxidation is brought 

about through the action of a microor- 
ganism present in the so-called mother of 
vinegar. The oxidation of alcohol into 
acetic acid through the agency of the 
organism is known as acetic fermenta- 
tion and may be represented as follows : 




C 2 H 5 OH 



0, 



H 



C 2 H 3 2 + H 2 



Manufacture of vinegar. The old 

method of making vinegar consisted 
simply in storing cider in barrels un- 
til the fermentation was complete — a 
process requiring several weeks. In the 
modern method the change is brought 
about in a few hours, a large cask known 
as a generator being used (Fig. 158). 
This is filled loosely with beech shav- 
ings. Vinegar is first sprayed into the 
top of the cask in order to introduce 
the mother of vinegar. The organism 
present attaches itself to the shavings, 
which are used because they present a 
large surface. Next a dilute solution of alcohol (hard cider, in 
the case of cider vinegar) is sprayed into the top of the cask 
while air is admitted at the bottom A, A. In this way the alcohol 
and oxygen are brought into intimate contact, and the oxida- 
tion takes place rapidly as the liquid trickles down over the shav- 
ings. The resulting vinegar is drawn off at the bottom B of the 
cask. Instead of starting with cider, one may use almost any sub- 
stance which contains starch or sugar, these compounds first being 
changed into alcohol, as explained in the manufacture of alcohol. 
In this way are prepared malt vinegar from starch and sugar 



Fig. 158. A generator for 
the manufacture of vinegar 



ORGANIC ACIDS AND THEIR DERIVATIVES 261 



vinegar from sugar residues. The cheapest vinegar is made from 
pure dilute alcohol and is known as distilled vinegar. It is colorless 
and leaves no residue upon evaporation. 

The Federal law requires that all vinegar shall contain not less 
than 4 per cent of acetic acid. In addition to the acid, vinegar pr3- 
pared from fruits and grains contains certain solids derived from 
the source materials. It is by studying the character of these 
solids left upon evaporating a sample of vinegar that the chemist 
is able to determine the source of the vinegar. 

The question might arise as to why it is not possible to make 
vinegar by simply adding to water the necessary percentage of 
acetic acid obtained from the distillation of wood. Such a product 
could be made at low cost and would be just as sour as vinegar ; 
however, it would not have the odor or flavor of vinegar. This is 
due to the fact that when acetic fermentation takes place, there 
are formed, in addition to acetic acid, small percentages of other 
compounds, and it is these which impart to 
vinegar its characteristic flavor and odor. 

Acids belonging to other series. In addi- 
tion to the fatty acids, the following deserve 
special mention : 

Oxalic acid (H 2 C 2 4 • 2 H 2 0). This is a white 
solid and is found in a number of plants. 

Malic acid (H 2 • C 4 H 4 5 • H 2 0). This acid 
occurs in apples, pears, and some other fruits. 

Tartaric acid (H 2 • C 4 H 4 6 ). This is a white 
solid and occurs in many fruits either in the 
free state or in the form of its salts. The 
acid potassium salt (KH . C 4 H 4 6 ) occurs in 
the juice of grapes. When the juice ferments 
in the manufacture of wine, this salt, being 
insoluble in alcohol, separates on the sides 
of the cask, and in this form is known as argol (Fig. 159). 
When purified it forms a white solid, which is sold under 
the name of cream of tartar and is used in baking-powders. 
The acid itself is often used in soft drinks. 




Fig. 159. Argols, 
the source material 
of cream of tartar 



2G2 



CHEMISTRY AND ITS USES 



Citric acid (H 3 • C 6 H 5 7 ). This acid occurs in the so-called 
citrous fruits, such as the lemon and grapefruit. It is a white 
solid, soluble in water, and is often used as a substitute for 
lemons in making lemonade. 

Oleic acid (H • C 18 H 33 2 ). The derivatives of this acid con- 
stitute the principal part of many oils and liquid fats. The 
acid itself is an oily liquid. 

Fats and oils. The hydrogen of acids can be replaced not only 
by metals but by hydrocarbon radicals (such as CH 3 and C 2 H 5 ) 
as well. The resulting compounds are termed esters. Some of 

these esters have a pleas- 
ant fruity odor and are 
used as flavoring agents. 
The main constituents 
of the common fats and 
oils, such as butter, 
lard, and olive oil, are 
esters of oleic, palmitic, 
and stearic acids and 
are known respectively 
as olein, palmitin, and 
stearin (Fig. 160). The 
radical present in these 
esters is C 3 H 5 . It is 
tervalent and is known as the glyceryl radical, since it is pres- 
ent in glycerin (C 3 H 5 (OH) 3 ). Since the glyceryl radical is 
tervalent, and since oleic, palmitic, and stearic acids are all 
monobasic, it is evident that three molecules of each acid must 
enter into the formation of each molecule of the ester derived 
from it. The relation in composition between these acids and 
the corresponding esters is shown in the following formulas : 

H • C 18 H 33 2 (oleic acid) C 3 H 5 (C 18 H 3 30 2 ) 3 (olein) 

H • C 16 H 31 2 (palmitic acid) .... C 3 H 5 (C 16 H 31 2 ) 3 (palmitin) 
11 • C 18 H 35 2 (stearic acid) C 3 H c( C i8 H 35°2)3 (stearin) 




Fig. 160. 



The three chief constituents of 
oils and fats 






ORGANIC ACIDS AND THEIR DERIVATIVES 263 

Olein is a liquid and is the main constituent of oils such as 
olive oil. Palmitin and stearin are white solids and are the 
chief constituents of the solid fats. Beef suet is principally 
stearin. 

The oils and fats are widely distributed in the animal and vege- 
table kingdoms. Almost every seed and fruit as well as the tissues 
of animals contain oil which can be removed by pressure. Thus 
we have olive oil, corn oil, cottonseed oil, coconut oil, palm oil, 
lard oil, fish oil, and a host of others. In separating these from 
the source material we are sure to obtain along with the oil some 
other substances dissolved in the oil. Sometimes these substances 
add to the flavor, as is the case with olive oil. More often they 
impair its looks and taste. Chemists have succeeded in finding 
ways for removing these impurities, and the oils when so refined 
are very much alike in composition. Thus the refined cottonseed 
oil is now largely used for the table in place of the more expen- 
sive olive oil. 

Changing oils into solid fats : the hydrogenation of oils. It 
will be noted that stearin differs from olein in composition in 
that it contains six more atoms of hydrogen in each molecule. 
Now if hydrogen is brought in contact with olein under proper 
conditions and hi the presence of a suitable catalytic agent 
(finely divided nickel is used), the olein takes up the additional 
hydrogen and is changed into the solid stearin. It is possible 
in this way to change the oils into solid fats. Certain commer- 
cial fats used in cooking, mostly sold under trade names, are 
made by this process from the comparatively inexpensive 
cottonseed oil. 

The cottonseed-oil industry. Cotton seeds are bulky, and formerly 
it was a source of expense to dispose of them. One of the functions 
of chemists is to find uses for apparently useless substances, and 
they soon discovered that cotton seeds # contain a valuable oil. In 
place of being a source of expense, cotton seeds are now the source 
of over one billion pounds of oil annually, worth many millions 
of dollars. 




A cotton plantation in 
the South, as shown in A. 
The cotton is picked by 
hand and placed in large 
baskets 



After picking, the cotton 
is hauled to the cotton gin 
as shown in B. This is a 
machine for separating 
the seeds from the cotton 



In C the seeds are care- 
fully separated from all 
foreign matter, such as 
sticks and pebbles, and 
then taken to the store- 
house 



Fig. 161, A, B, and C. The cottonseed oil industry 



In D the clean seeds are 
crushed and heated, then 
placed in large presses 
and subjected to great 
pressure. In this way the 
oil is squeezed out. The 
solid matter left (about 
20 per cent by weight of 
the seeds) is a valuable 
cattle food 



The crude oil obtained in 
D is dark in color and 
is refined by pouring it 
into large tanks (E) and 
treating it with sodium 
hydroxide and other com- 
pounds and finally with 
a porous clay known as 
fuller's earth 



In F are shown in jar 1 the 
crude cottonseed oil just 
as it comes from the seed, 
in jar 2 the refined oil, and 
in jar 3 the oil after its 
treatment with hydrogen 
(hydrogenated oil) 









Fig. 161, D, E, and F. The cottonseed oil industry 



264 CHEMISTRY AND ITS USES 

Butter fat and oleomargarine. While butter fat, like other 
fats, consists principally of olein, palmitin, and stearin, its 
characteristic flavor is due to the presence of a small amount 
(about 8 per cent) of the fat butyrin, which is an ester of 
butyric acid and has the formula C 3 H 5 (C 4 H 7 2 ) 3 . Oleomar- 
garine differs from butter mainly in the fact that a smaller 
amount of butyrin is present. It is made from the fats obtained 
from cattle and hogs. Sometimes cottonseed and coconut oils 
are also added or are used alone. These fats are churned with 
milk or mixed with a small amount of butter in order to 
furnish sufficient butyrin to give the butter flavor. 

In appearance oleomargarine differs from most butter in being 
nearly colorless. While it is a common practice to color butter 
artificially, the Federal law permits the coloring of oleomargarine 
only upon the payment of a tax of 10 cents for each pound colored. 
Many of the states, however, have laws forbidding the sale of 
oleomargarine that is artificially colored, even though the Federal 
tax has been paid. 

Proteins. The term protein is applied to a large class of 
complex nitrogenous compounds which are everywhere abun- 
dant in animal and vegetable organisms and which constitute 
the principal part of the tissues of the living cell. The casein 
of milk, gluten of flour, and albumin of egg will serve as exam- 
ples of typical protein matter. The proteins all contain nitrogen, 
carbon, hydrogen, and oxygen, and some contain sulfur and 
phosphorus in addition. 

EXERCISES 

1. Name the different kinds of fermentation studied and the changes 
involved in each. 

2. Why does it take so long for a barrel of cider to turn into 
vinegar? 

3. What changes are involved in making vinegar from starch? 



ORGANIC ACIDS AND THEIR DERIVATIVES 265 

4. Distilled vinegar is cheaper than cider vinegar, but it is colorless 
'or nearly so. Sometimes it is colored with caramel to imitate cider 

vinegar. How could you distinguish between such a vinegar and pure 
cider vinegar? 

5. What effect did the passage of the Prohibition Act have on the 
production of cream of tartar ? 

6. For what purpose is oleomargarine artificially colored? 

7. For what purpose is butter artificially colored ? 

8. Why does the Federal government permit the coloring of butter 
but not of oleomargarine unless a tax is paid ? 

9. Is it possible by the taste alone to distinguish (a) between refined 
cottonseed oil and olive oil ? (b) between butter and oleomargarine ? 

10. Give examples of substances that the chemist has changed from 
being a source of expense to a source of great wealth. 

11. Ammonium oxalate is a very common chemical reagent. What 
is its formula ? 

12. Magnesium citrate is a common medicine. What is its formula? 

13. Ordinary rhubarb used as a food contains the potassium acid 
salt of oxalic acid. What is its formula? 

14. Would cider change into vinegar if all air were excluded? 

15. The oils obtained from petroleum (such as coal oil and lubricat- 
ing oil) are often called mineral oils to distinguish them from the animal 
and vegetable oils described in this chapter. How do these two classes 
of oils differ in composition ? 

16. What weight of the alcohol of the druggist is necessary to make 
100 kg. of distilled vinegar (contains 4 g. of acetic acid in 100 g. of 
vinegar) ? 



CHAPTER XXIX 
FOODS 

Composition and function of foods. While the compounds 
present in our foods are very numerous and often exceed- 
ingly complex, yet they may all be included under a few 
general heads ; namely, proteins, fats, 'carbohydrates, mineral 
matter, and water. Since the mineral matter is left as a resi- 
due when the food is burned, it is listed as ash in reporting 
the analysis of foods. In addition to the above constituents 
many foods contain small percentages of substances known 
as vitamins, which are essential to life. The composition of 
the more common foods is given in the table on page 268. 

The function and uses of food are also clearly stated in 
the chart, on the opposite page, prepared by the United States 
Department of Agriculture. The facts included in this chart 
are of great importance and should be carefully studied. 

While the different classes of food materials are to a cer- 
tain extent interchangeable, experiments show that a proper 
mixture of these materials is essential to health. Of course 
it is true that one can live for many days on a purely protein 
diet or on a diet purely of fats and carbohydrates; in fact, 
persons have been known to live for many days without any 
food whatever (other than water). In all such cases the body 
derives the necessary materials from the surplus supply always 
stored up in the normal body. 

Vitamins. Experiments in recent years have shown that 
in order for the body to grow and maintain its health there 
must be present in our foods, in addition to protein, fats, 

266 



Chart IA-. Composition, Functions, And Uses Of Food. 



Revised Edition. 



US. Department of Agriculture Prepared by 

Office of Experiment Stations C. R LANGWORTHY 

A.CTrue. Director. Expert in charge of Nutrition Investigations 

FUNCTIONS AND USES OF FOOD. 
CONSTITUENTS OF FOOD. 



Water 



FOOD AS PUR- 
CHASED CONTAINS 



EDIBLE PORTION < 
Flesh of meat.yolk [ Nutrients 
and white of eggs, 
wheat flour, etc. 
REFUSE! 
Bones, entrails, 
shells, bran.etc. 

USE OF FOOD IN THE BODY. 



PROTEIN - Builds and repairs tissue 

White (albumen) of eggs, 

curd (casein) of milk, 

lean meat.gluten of wheat.etc. 



Protein 
Fats 

Carbohydrates 
Mineral Matter Or Ash 



All serve as fuel to 
yield energy inthe forms 
of heat and muscular 



power 



FATS Are stored as fat 

Fat of meat.butter, 

olive oil, oils of corn 

and wheat, etc. 
CARBOHYDRATES— -Are transformed into fat 

Sugar, starch, etc. 
MINERAL MATTER OR ASH— Share in forming bone, 

Phosphates of lime, assists in digestion, etc. 

potash, soda, etc. 

Food isthatwhich, taken into the body, builds tissue or yields energy. 



268 



CHEMISTRY AND ITS USES 



AVERAGE COMPOSITION OF EDIBLE PORTION OF TYPICAL 
FOODS EXPRESSED IN GRAMS PER 100 GRAMS OF FOOD 















Fuel 














Value 


Food 


Water 


Protein 


Fat 


HYDRATE 


Ash 


(Cal. per 
100 g.) 


Almonds 


4.8 


21.0 


54.9 


17.3 


2.0 


647 


Apples 


84.6 


0.4 


0.5 


14.2 


0.3 


63 


Asparagus 


94.0 


1.8 


0.2 


3.3 


0.7 


22 


Bacon (smoked) . . . 


20.2 


9.9 


64.8 





5.1 


623 


Bananas 


75.3 


1.3 


0.6 


22.0 


0.8 


99 


Beans (dried) .... 


12.6 


22.5 


1.8 


59.6 


3.5 


345 


Beans (string) .... 


89.2 


2.3 


0.3 


7.4 


0.8 


42 


Beef (lean steak) . . . 


70.0 


21.0 


7.9 





1.1 


155 


Beef (slightly fat) . . 


73.8 


22.1 


2.9 





1.2 


115 


Beets 


87.5 


1.6 


0.1 


9.7 


1.1 


46 


Bread (corn) 


38.9 


7.9 


4.7 


46.3 


2.2 


259 


Bread (graham) . . . 


35.7 


8.9 


1.8 


52.1 


1.5 


260 


Bread (white) .... 


35.3 


9.2 


1.3 


v 53.1 


1.1 


260 


Butter 


11.0 


1.0 


85.0 





3.0 


769 


Cabbage 


91.5 


1.6 


0.3 


5.6 


1.0 


32 


Carrots 


88.2 


1.1 


0.4 


9.3 


1.0 


45 


Celery 


94.5 


1.1 


0.1 


3.3 


1.0 


19 


Chestnuts 


45.0 


6.2 


5.4 


42.1 


1.3 


242 


Chicken 


63.7 


19.3 


16.3 





1.0 


224 


Codfish (fresh) .... 


82.6 


15.8 


0.4 





1.2 


67 


Corn (green) 


75.4 


3.1 


1.1 


19.7 


0.7 


101 


Dates 


13.8 


1.9 


2.5 


70.6 


1.2 


313 


Eggs 


73.7 


14.8 


10.5 





1.0 


154 


Figs . 


18.8 


4.3 


0.3 


74.2 


2.4 


317 


Ham (lean, smoked) . . 


53.5 


20.2 


20.8 





5.5 


268 


Lettuce 


94.7 


1.2 


0.3 


2.0 


0.9 


16 


Macaroni 


78.4 


3.0 


1.5 


15.8 


1.3 


89 


Milk 


87.0 


3.3 


4.0 


5.0 


0.7 


69 


Oatmeal 


7.3 


16.1 


7.2 


67.5 


1.9 


400 


Olive oil 








100.0 








900 


Oranges 


86.9 


0.8 


0.2 


11.6 


0.5 


51 


Peaches 


89.4 


0.7 


0.1 


9.4 


0.4 


41 


Peanuts 


9.2 


25.8 


38.6 


24.4 


2.0 


548 


Peas (green) 


74.6 


7.0 


0.5 


16.9 


1.0 


100 


Plums 


78.4 


1.0 





20.1 


0.5 


84 


Potatoes 


78.3 


2.2 


0.1 


18.4 


1.0 


83 


Prunes (dried) .... 


22.3 


2.1 





73.3 


2.3 


302 


Baisins 


14.6 


2.6 


3.3 


76.1 


3.4 


345 


Rice 


12.3 


8.0 


0.3 


79.0 


0.4 


351 


Salmon 


64.6 


21.2 


12.8 





1.4 


200 


Spinach 


92.3 


2.1 


0.3 


3.2 


2.1 


24 


Strawberries .... 


90.4 


1.0 


0.6 


7.4 


0.6 


39 


Tomatoes 


94.3 


0.9 


0.4 


3.9 


0.5 


23 


Turnips 


89.6 


1.3 


0.2 


8.1 


0.8 


40 


Wheat flour 


11.9 


13.3 


1.5 


72.7 


0.6 


357 



These values are taken from Bulletin No. 28, office of Experiment Station, Washing- 
ton, D. C. The fuel values are obtained from the following formula : 

Cal. in 100 g. =4 P + 9 F+ 4 C in which P, F, and C represent respectively the num- 
ber of grams of protein, fat, and carbohydrates in 100 g. of the food. 



FOODS 269 

carbohydrates, and mineral matter, small amounts of certain 
substances which are known as vitamins. But little is known 
of them beyond their physiological effects. Foods from which 
they have been extracted, although containing the necessary 
amounts of protein, carbohydrates, and fats, no longer nourish 
the body and maintain health. At least three kinds of vita- 
mins are known to exist. The names of these, together with 
the function of each and some food which is particularly rich 
in each, are as follows : 

1. Fat soluble A. Promotes growth, keeps the body in good 
condition, and thus prevents disease in general. Present in 
milk and butter. 

2. Water soluble B. Prevents the disease known as beri-beri. 
Present in fresh vegetables and yeast. 

3. Water soluble C. Prevents scurvy. Present in the juices 
of the tomato, the orange, and the lemon. 

All three kinds are abundant in green vegetables such as 
lettuce and spinach and in milk. Just how these vitamins 
act to promote growth and prevent disease is not known. 
It may be that their function is something like that of a 
catalytic agent, but this is simply a guess. 

Energy value of foods. Experiments show that the heat 
of the body, as well as the energy used in muscular work, 
results from the oxidation of food materials. The foods, 
when eaten, undergo complex changes in which the insoluble 
portions are converted into soluble compounds. These are 
absorbed into the system and then either undergo oxidation 
directly or are temporarily built into tissues which later un- 
dergo oxidation. In this process most of the carbon is finally 
changed into carbon dioxide and exhaled from the lungs, 
while the hydrogen is changed into water. The nitrogen is 
excreted largely in the form of urea (CO(NH 2 ) 2 ). 

Broadly speaking, foods may be regarded as fuel from 
the oxidation of which in the body the energy necessary for 



270 



CHEMISTRY AND ITS USES 



physical requirements is set free. In the study of foods it is 
convenient, therefore, to use their fuel values (heats of com- 
bustion) as a basis of comparison. These values are deter- 
mined in the calorimeter and expressed in large calories (Cal.), 
which are 1000 times as large as the small calorie (cal.). 

Now experiments show that the body obtains from each of 
the three classes of foods, when absorbed in the system and 
oxidized, approximately the following fuel values: 



Class of Foods 



Carbohydrates 
Fats . . . . 
Proteins 



Calories 
per Gram 



Calories 
per Pound 



1815 

4082 
1815 



Amount and nature of foods necessary for health. Many 
studies have been made in order to find out just how much 
and what kind of food is best adapted for the preservation of 
health. Evidently many conditions, such as one's age, weight, 
and occupation and the climate in which one lives, enter into 
the problem. 

Since the fats and carbohydrates have nearly the same func- 
tion, it is sufficient in stating food requirements for a period 
of, say, twenty-four hours, to give simply the weight of pro- 
tein necessary, together with the total fuel value. The differ- 
ence between the total fuel value and that of the required 
protein gives the number of calories to be supplied from fats 
and carbohydrates. Such a mixture of these two food mate- 
rials is selected so that the fuel value of the mixture, together 
with the fuel value of the protein, equals the total fuel value 
required. Different dietary standards have been proposed, 
but the following (p. 271) given by Atwater are generally 
accepted. They give the food requirements for a period of 
twenty -four hours. 






FOODS 



271 



Character of Individual 



Proteins 


Fuel Value 


required 


required 


175 g. 


5500 Cal. 


125 g. 


3400 Cal. 


112 g. 


3050 Cal. 


100 g. 


2700 Cal. 


90 g. 


2450 Cal. 


108 g. 


3060 Cal. 


100 g. 


2700 Cal. 


87 g. 


'2380 Cal. 


75 g. 


2040 Cal. 


62 g. 


1700 Cal. 



Man with very hard muscular work (wood- 
chopper, football player) 

Man with moderately active muscular work . 

Man with light to moderate muscular work . 

Man at sedentary occupation or woman with 
moderately active work 

Man at rest or woman with light muscular 
work 

Boy 15 to 16 years 

Boy 13 to 14 years or girl 15 to 16 . . 

Boy 12 to 13 years or girl 14 to 15 . . . . 

Boy 10 to 11 years or girl 10 to 12 . . . . 

Boy 6 to 9 years . 



In selecting onr foods, however, it is essential not only that 
we secure the proper proportion of protein, fats, carbohydrates, 
and mineral matter, but we must also consider whether the 
foods selected contain these ingredients in a palatable and 
digestible condition. Especially must we make such a selec- 
tion as will furnish us with the necessary vitamins ; in other 
words, we must take into consideration quality as well as quan- 
tity of food. For these reasons some foods, such as milk, may 
not seem to rank high as an economical food considered 
simply from its content of fat, carbohydrate, and protein ; 
yet when we take into account the character of these con- 
stituents present in milk and especially its vitamin and 
mineral content, its effect in maintaining health as well as in 
nourishing the body makes it one of our most valuable foods. 

Laws of the conservation of energy and matter in the human body. 

In studying the food requirements of the human body many ex- 
periments have been made to find out whether or not the laws 
holding good in the inanimate world (namely, the laws of the con- 
servation of matter and energy) hold good also in the processes 
taking place in the living organism. So far as the law of 



272 CHEMISTRY AND ITS USES 

conservation of matter is concerned, the question can be answered 
by carefully comparing the air breathed and the food eaten 
(intake) with the products stored up in the body and with the 
products exhaled or excreted (output). 

In order to measure the energy changes, however, one must be 
able to determine the amount of heat evolved. This is done by 
means of the respiration calorimeter. This is a chamber large 




Fig. 162. Exterior of a respiration calorimeter 

enough to enable one to live in it and constructed with double 
walls of nonconducting material so as to prevent loss of heat 
through ladiation; in fact, it is a large calorimeter (Fig. 126). 

With the necessary precautions it is possible in this way to 
measure the heat generated in the body of a person living within 
the calorimeter. The experiments show that the changes in matter 
and energy which take place in the human body are in accord 
with the laws of the conservation of matter and energy. 

Fig. 162 shows the exterior of a respiration calorimeter, also 
the apparatus for analyzing the air admitted and measuring the 
heat evolved by the person in the calorimeter. 



FOODS 



273 



Cost of food as related to its nutritive value. It will 
be noted that the food requirements for the body are stated 
simply in terms of protein, carbohydrates, and fats, but such 
a selection of these must be made as will contain the neces- 
sary vitamins. It is very evident that the cost of, say 100 g., 
of protein will vary according to the source of the protein. 

wsmm 





Fig. 163. The label on a well-known patent medicine before and after the 
passage of the Pure Food and Drugs Act 

For example, protein obtained in the form of tenderloin steak 
will cost much more than an equal weight of protein obtained 
from dried beans. Again, there is as much nutriment in 1 lb. 
of wheat flour as in 3J qt. of oysters, although the latter is 
far more expensive. By finding the selling prices of the 
various foods given in the table on page 268, one can easily 
determine the relative cost of various food materials obtained 
from different sources. 



274 CHEMISTRY AND ITS USES 

Pure Food and Drugs Act. In 1906 the Federal congress passed 
wliat is known as the Pure Food and Drugs Act. This fixes the 
standard of various foods and drugs, and prescribes that all labels 
must tell the truth and must state the presence of any dangerous 
or habit-forming drugs. The effect of the act may be seen by a 
study of Fig. 163, which shows the label used on a well-known 
patent medicine before the act was passed and also how it had to 
be modified to meet the requirements of the act. 

EXERCISES 

1. Is the ash obtained in burning a food necessarily present in that 
form ? 

2. Why do different nations use different kinds of foods ? 

3. Why do we enjoy certain kinds of foods more in winter than in 
summer ? 

4. Mention foods that are rich in water ; in protein ; in fats ; in 
carbohydrates. 

5. From the current prices of foods work out a list of the most eco- 
nomical ones to use as a source of protein ; of carbohydrates ; of fats. 

6. If one wishes to grow thin, what kinds of foods ought to be 
avoided ? 

7. In selecting a proper diet what factors must we take into 
consideration ? 



CHAPTER XXX 
THE PHOSPHORUS FAMILY 



Name 

Phosphorus . . . 

Arsenic . . . . 

Antimony . . . 

Bismuth . . . , 



Symbol 



Atomic 
Weight 



Density 



Melting 

POINT 



P 

As 
Sb 
Bi 



31.04 
74.96 
120.2 

208. 



1.83 
5.73 
6.70 
9.80 



630 c 
271 c 



The family. The elements constituting this family belong 
in the group with nitrogen and therefore resemble it in a 
general way. They exhibit a regular gradation of properties, 
as is shown in the above table. 



Phosphorus 

Properties. We commonly associate the element phosphorus 
with matches and with bones, for it is well known that it is 
the substance used to make matches inflammable and that its 
compounds constitute the principal mineral constituents of 
bones. The free element is obtained in two chief forms, 
known as white (or yellow) phosphorus and red phosphorus. 

White phosphorus (yelloiv phosphoi-us) is a nearly colorless, 
translucent, waxy solid which melts at 44°. It is sold in the 
form of sticks (Fig. 164) and is the common form of phosphorus. 
Red phosphorus, on the other hand, is a dark-red powder. 

The two allotropic forms of phosphorus differ widely in proper- 
ties. Thus the white form is very poisonous, is soluble in carbon 
disulfide, and ignites so readily that it must be kept under water 

275 



270 



CHEMISTRY AND ITS USES 



and handled with the greatest care. It is easily cut into pieces, 
but this must always be done under water to avoid ignition. Red 
phosphorus, on the other hand, does not ignite easily, is not solu- 
ble in carbon disulfide, is not poisonous, and has a high melting 
point. White phosphorus is formed when any form of the element 
is distilled and its vapor condensed in cold water. On standing 
at ordinary temperatures it slowly changes into the red variety ; 
at higher temperatures the change is much more rapid. 




Fig. 164. Phosphorus and some of its most common compounds 

History and occurrence. Phosphorus has long been known, 
having been discovered in 1669 by the alchemist Brand while 
searching for the philosopher's stone. It occurs almost entirely 
in various mineral forms of calcium phosphate (Ca 3 (P0 4 ) 2 ). 
Phosphorite is the most abundant of these minerals, while 
apatite consists of calcium phosphate, together with calcium 
fluoride or chloride. Calcium phosphate is the chief mineral 
constituent of the bones of animals, and bone ash is therefore 
nearly pure calcium phosphate. 

Preparation. The element is prepared by heating a mix- 
ture of calcium phosphate, sand, and carbon in an electric 
furnace. Phosphorus is liberated as a vapor, and this is 
condensed under cold water and cast into stick form as the 
white variety. 






THE PHOSPHORUS FAMILY 277 

Chemical conduct. The most striking property of white 
phosphorus is its affinity for oxygen. When exposed to cool 
air it slowly combines with oxygen, and in so doing gives out 
a pale light, or phosphorescence, which can be seen only in a 
dark place ; hence the word phosphorus, which means " light 
producer." The heat of the room may raise the temperature 
of phosphorus to the kindling point, when it burns with a 
sputtering flame, giving off dense fumes of oxide of phos- 
phorus. It burns with dazzling brilliancy in oxygen and com- 
bines directly with many other elements. On account of its 
great attraction for oxygen it is preserved under water. 

Compounds. While phosphorus forms many compounds, 
only a few of them are of common importance. 

1. Oxides. Phosphorus forms two well-known oxides — 
the trioxide (P 2 3 ) and the pentoxide (P 2 O g ), sometimes 
called phosphoric anhydride. The pentoxide is much the better 
known of the two. It is formed when phosphorus burns 
in oxygen or air and is a snow-white, voluminous powder 
(Fig. 164) whose most marked property is its great attraction 
for water. It has no chemical action upon most gases, so that 
they can be very thoroughly dried by being passed through 
properly arranged vessels containing phosphorus pentoxide. 

2. Sulfide. When phosphorus and sulfur are brought together 
under certain conditions, they combine to form phosphorus 
sesquisulfide (P 4 S 3 ). This is a dark-colored solid and is largely 
used in making matches. 

3. Acids. Phosphorus forms a number of acids, but the only 
very important one is phosphoric acid (H 3 P0 4 ). The pure com- 
pound forms large, colorless crystals ; ordinarily the acid con- 
tains a little water, and it then forms a thick, sirupy liquid 
(Fig. 164). If a solution of the pure acid is desired it is pre- 
pared by burning phosphorus and adding water to the pentoxide 
so formed: 

3H 2 + P 2 5 -^2H 3 P0 4 



278 CHEMISTRY AND ITS USES 

The ordinary solution, however, is prepared by treating calcium 
phosphate with sulfuric acid : 

Ca 3 (P0 4 ) 2 + 3 H 2 SO — ► 3 CaS0 4 + 2 H 3 PO, 

The calcium sulfate (CaS0 4 ) formed is nearly insoluble and 
can be separated from the phosphoric acid by nitration. Being 
a tribasic acid, it forms acid salts as well as normal salts. Thus 
the following sodium salts are known : 

NaH 2 P0 4 primary, or monosodium-hydrogen phosphate 

Xa 2 HP0 4 secondary, or disodium-hydrogen phosphate 

Na 3 P0 4 tertiary, or normal sodium phosphate 

Uses. Because of its poisonous properties white phosphorus 
is used as a rat poison. It is also used in making pure phos- 
phoric acid and for minor purposes in chemical laboratories. 
During the World War large amounts were used in generating 
smoke clouds (P 2 5 ) for concealing troops and ships (Fig. 165). 
The incendiary bullets which proved so disastrous to the 
Zeppelins during the war contained free phosphorus. When 
fired, the phosphorus was ignited by the friction of the air, 
and, on hitting a balloon or Zeppelin, set fire to the hydrogen 
with which they were filled. The chief use of phosphorus, 
however, is in the making of matches. 

Matches. Friction matches containing phosphorus first came 
into use in 1827, and at present two general varieties are in common 
use. The more common variety which will ignite when rubbed 
against any rough surface is made by dipping the match stick 
first into some inflammable substance, such as melted paraffin, and 
afterwards into a paste consisting of (1) phosphorus sesquisulfide, 
P 4 S 8 ; (2) some oxidizing substance, such as manganese dioxide, 
red lead, or potassium chlorate ; and (3) a binding material, such' 
as glue or dextrin. On friction the phosphorus is ignited, the com- 
bustion being supported by the oxidizing agent and communicated 
to the wood by the burning paraffin. In sulfur matches the paraffin 
is replaced by sulfur, 



THE PHOSPHORUS FAMILY 



279 



In the safety match red phosphorus, an oxidizing agent, and 
some gritty material, such as powdered glass, are mixed with 
glue and placed on the side of the box. The match tip is 
provided with an oxidizing agent and an easily combustible 
substance, usually antimony sulfide. The match cannot be 
ignited easily by friction except on the prepared surface. 




Fig. 165. The use of phosphorus in the World War for concealing troops 

Often when troops were ready for an attack, shells filled with phosphorus were 

fired. When these exploded, the phosphorus burned, forming dense clouds of the 

oxide which entirely concealed the troops 



Constant working with white phosphorus frequently results in 
dreadful diseases of the bones of the face, while many disastrous 
fires are caused by accidental ignition of matches containing it. 
On this account the manufacture and use of such matches is 
prohibited by law in many countries. The congress of the United 
States in 1913 accomplished the same end by imposing a prohibitive 
tax upon white-phosphorus matches. 

Phosphates. The phosphates form an important class of salts. 
The most important phosphate is calcium phosphate, which is 
mined in enormous quantities for use as a fertilizer. 



280 



CHEMISTRY AND ITS USES 



Arsenic 

General discussion. When we hear the word arsenic we 
are apt to think of a poisonous substance because it is well 
known that materials containing arsenic are poisonous in 
character. The element itself is a steel-gray solid and occurs 
in nature combined with sulfur. Small amounts of the ele- 
ment combined with metals are associated with copper ores. In 

the separation of copper from these 
ores the arsenic is obtained as a by- 
product in the form of the oxide 
As 2 3 . The element can be obtained 
from this by heating with carbon : 




2 As O + 3 C 



As + 3 CO, 



Fig. 166. The two common 
forms of arsenic trioxide 



Arsenic burns when heated in air, 
forming the oxide As 2 3 . 

Compounds. When hydrogen is 
generated in contact with arsenic or 
its compounds a very poisonous, color- 
less gas known as arsine (AsH 3 ) is 
generated. The most common com- 
pound of arsenic is the oxide As 2 3 , 
known as arsenic trioxide or arsenious oxide, or often as white 
arsenic. It is a rather heavy white solid, obtained either as a 
crystalline powder or in lumps resembling porcelain in appear- 
ance (Fig. 166). It is very poisonous, from 0.2 to 0.3 g. being 
a fatal dose. Arsenic also forms a number of acids, the most 
important of which are arsenious acid (H g As0 3 ) and arsenic 
acid (H 3 As0 4 ). Some of the salts of these acids are of con- 
siderable importance. 

Uses. Arsenic in the free state is little used. About 0.5 per 
cent of arsenic is added to lead used in making shot, since 
this makes the lead harder. The compound of arsenic most 






THE PHOSPHOEUS FAMILY 281 

widely used is the oxide As 2 O g . This is used in preparing 
all the other compounds of arsenic. It is also used in making 
glass and as a preservative in mounting the skins of birds and 
animals. Its greatest use is in the preparation of insecticides 
for spraying trees and killing insects on stock. 

Arsenic insecticides. Several compounds of arsenic have 
important uses as insecticides. Paris green and Scheeles 
green are complex salts made by treating solutions of cop- 
per salts with arsenious oxide. The commercial lead arse- 
nate so widely used in connection with lime-sulfur sprays 
(p. 185) is chiefly Pb 3 (As0 4 ) 2 . 

Antimony 

General discussion. Antimony is a silverlike, brittle solid 
a little lighter than iron. It melts at 630° and expands some- 
what on solidifying. It is found in nature chiefly in the form of 
stibnite (Sb 2 S 3 ), from which it is obtained by heating with iron: 

Sb 2 S 3 + 3 Fe — h 2 Sb + 3 FeS 

It forms a number of compounds which resemble the cor- 
responding compounds of arsenic in composition and prop- 
erties. Its chief use is in making alloys (p. 282), especially 
the one known as Babbitt metal. 

Bismuth 

General. Bismuth resembles antimony in appearance except 
that it has a slightly rosy tint. It is a little heavier than iron, 
but melts at a low temperature (271°). It occurs in nature 
chiefly in the free state, and our supply comes largely from 
Bolivia. Like antimony, its chief use is in the making of 
alloys. The element differs from the other elements of the 
phosphorus family in that it has practically no acid proper- 
ties, but forms salts like the metals. Sulfuric and nitric 



282 CHEMISTRY AND ITS USES 

acids dissolve it just as they do copper. When heated in 
the air it burns, forming the oxide Bi 2 8 , and this dissolves 
in acids, forming salts: 

Bi 2 3 + 6 HC1 ^->- 2 BiCl 8 + 3 H 2 
Bi 2 3 + 6 HN0 3 — » 2 Bi(N0 3 ) 3 + 3 H 2 

These salts are colorless, crystalline solids. 

Action of water on bismuth salts. The salts of bismuth react 
with water as shown in the following equations: 

BiCl 3 + H 2 q=± BiOCl + 2 HC1 
Bi ( N0 3 ) 3 + H 2 ° =t=± BiONO, + 2 HN0 3 

The compound BiOCl, known as bismuth oxychloride, as well 
as the corresponding bismuth oxynitrate, BiON0 3 , are white 
solids, the latter being often used in medicine. 

Alloys 

Bismuth and antimony and many of the metals when melted 
together thoroughly intermix, and on cooling form a uniform 
metal-like substance called an alloy. Not all metals will mix 
in this way, and in some cases definite chemical compounds 
are formed and separate out as the mixture solidifies, thus 
destroying the uniform quality of the alloy. In general, the 
melting point of the alloy is below the average of the melting 
points of its constituents. 

Both antimony and bismuth alloy readily with many of the 
metals. The alloys so formed are heavy, are easily melted, 
do not oxidize easily nor act upon water, and, in general, are 
well adapted to many technical uses. The manufacture of 
alloys constitutes the chief use of these two elements. 

Antimony imparts to its alloys the property of expanding 
slightly in solidification, which renders them especially use- 
ful in type founding, where fine lines are to be reproduced on 



THE PHOSPHOEUS FAMILY 



283 



a cast. Type metal consists of antimony, lead, and tin. Babbitt 
metal, largely used in the construction of certain parts of 
machinery, contains the same metals in a different proportion, 
together with a small percentage of copper. 

Bismuth is particularly valuable in the production of very 
low-melting alloys. For example, Wood's metal, consisting of 
bismuth, lead, tin, and cadmium, melts at 60.5°. The low 
melting point of such alloys is turned to practical account 



|jS 








% 








■ ' 1 1 


!:. . . ..._ " 
















4; . ; : ; ., „ 







Fig. 167. An automatic fire curtain above a door 

in making safety plugs in boilers, automatic fire curtains 
and automatic water sprinklers in buildings, and in many 
similar devices. 

Fig. 167 shows a fire curtain, which is held in place by two 
wires A, A, joined at B by a bismuth alloy. In case of fire the 
alloy melts and the wires holding the curtain up are thereby 
released and the curtain drops, covering the door. 



EXERCISES 

1. What elements have we studied that occur in allotropic forms? 

2. What substances have we studied that are used as insecticides? 

3. Nitrogen occurs in the same group in the periodic table as the 
phosphorus family. Can you point out any similarity in chemical 
conduct ? 

4. Approximately, what weight of phosphorus is present in your 
body (p. 10) ? In what form is it present ? 



284 CHEMISTRY AND ITS USES 

5. By far the greater part of the phosphorus used in the United 
States is made at Niagara Falls. What is the advantage of this locality 
for the production of phosphorus ? 

6. In what connection has phosphorus been referred to in an 
earlier chapter ? 

7. How does an automatic water sprinkler in a building operate to 
extinguish fires? 

8. In the equation on page 280 what does the character As 4 indicate ? 

9. Write the formula and give the name (a) of the sodium salt of 
arsenious acid ; (b) of arsenic acid. 

10. Phosphorus pentoxide combines with water to form phosphoric 
acid (p. 277). To what class of compounds does this oxide belong ? 

11. Why does cutting a piece of phosphorus cause it to ignite? 

12. What is the function of the potassium chlorate used in ordinary 
matches ? 

13. 100 kg. of pure calcium phosphate would yield what weight of 
phosphorus ? 

14. What weight of phosphoric acid could be prepared by burning 
phosphorus in air and adding water to the product ? (Suggestion. 1 atom 
of phosphorus forms 1 molecule of phosphoric acid.) 

15. What weight of antimony could be prepared from 1 ton of 
stibnite ? 



CHAPTER XXXI 
SILICON AND BORON 

Silicon 

General. Just as carbon is the most essential element pres- 
ent in living matter (plants and animals), so silicon is the 
most essential element present in the compounds that con- 
stitute the soils and the rocks that make up the earth's crust. 
It is therefore very abundant, being next to oxygen in this 
respect. Notwithstanding this fact it is not so familiar as 
many other elements of rarer occurrence. This is because it 
never occurs native. It can be prepared by heating a mixture 
of sand (Si0 2 ) and carbon in an electric furnace : 

2 C + Si0 2 — y Si + 2 CO 

It is a bright, metallic, brittle solid and is used largely for 
making certain varieties of iron and iron alloys. 

Silicon occurs in nature chiefly in the form of silicon diox- 
ide (silica), Si0 2 , or in the form of salts of silicic acids (sili- 
cates). These compounds constitute almost all the common 
rocks except limestone. Plants absorb small amounts of silica 
from the soil, and it is also found in minute quantities in 
animal organisms, especially in hair, claws, and horns. 

Silicides ; carborundum. As the name indicates, silicides 
are compounds consisting of silicon and some one other ele- 
ment. They are very stable at high temperatures and are 
usually made by heating the appropriate substances in an 
electric furnace. 

285 



28G 



CHEMISTRY AND ITS USES 



The most important silicide is carborundum, which is a 
silicide of carbon of the formula CSi. When carbon is heated 
with sand in an electric furnace, the element is - first liberated 
as shown under the preparation of silicon. Under proper con- 
ditions the silicon then combines with excess of carbon to 
form CSi, and it is in this way that carborundum is prepared. 







m 



Fig. 168. Crystals of carborundum and abrasive utensils made 
of carborundum 



Large quantities of this product are made at Niagara Falls. 
It can be obtained as perfectly colorless crystals, but as made 
commercially the crystals have a beautiful purplish-black 
color, the color being imparted to it by the presence of small 
amounts of certain other elements, chiefly iron. Carborundum 
is very hard, being surpassed in this property only by the 
diamond and one or two rare compounds. It is this property 
which makes it so valuable. It is used as an abrasive ; that is, 
as a material for grinding and polishng very hard substances. 
Fig. 168 shows two samples of the crystalline material, as 
well as whetstones and a grinding-wheel prepared from 
carborundum. 



SILICON AND BORON 287 

Manufacture of carborundum. The materials used in making 
carborundum are carbon, sand, common salt, and sawdust. The 
salt assists in the reaction, while the sawdust burns away, leav- 
ing the mass porous and thus allowing the escape of gases. The 
mixture of these materials is heated in a large resistance fur- 
nace similar to the one employed in the manufacture of graph- 
ite (p. 104). Fig. 169 shows a furnace in operation. The carbon 
monoxide formed escapes and burns as shown in the figure. 




Fig. 169. A carborundum furnace in operation 

Silicon dioxide (silica) (Si0 2 ). This substance is found in a 
great variety of forms in nature, both in the amorphous and 
in the crystalline condition. In the form of quartz (Fig. 170) 
it is found in beautifully formed six-sided prisms, sometimes 
of great size. When pure it is perfectly transparent and color- 
less. Some colored varieties are given special names, as ainethyst 
(violet), rose quartz (pale pink), smoky or milky quartz (colored 
and opaque). Other varieties of silicon dioxide, some of which 
also contain water, are chalcedony, onyx, jasper, opal, agate, 
and flint. Sand and sandstone are largely silicon dioxide. 

The skeletons of certain microorganisms (infusoria) are 
composed of nearly pure silica. In some localities these have 



288 



CHEMTSTKY AND ITS USES 



accumulated in immense deposits, forming a very fine and 
sharp sand called infusorial earth. This material is often used 
as a scouring-powder, especially in scouring-soaps. 

Properties. As obtained by chemical processes, silicon diox- 
ide is an amorphous white powder. In the crystallized state it 
is very hard and has a density of 2.6. Pure silica begins to 
soften at about 1600°, and somewhat above this temperature 

it can be drawn out 
into threads and blown 
like glass into tubes and 
small vessels. These 
articles are attacked by 
comparatively few or- 
dinary reagents and do 
not expand or contract 
to any appreciable ex- 
tent with even very 
great changes in tem- 
perature. On this ac- 
count a quartz vessel 
can be heated red hot 
and plunged into cold 
water without cracking. Unfortunately they are quite expen- 
sive, but their use is steadily increasing. Fig. 171 shows a 
quartz crucible and quartz tubes on a wire triangle used 
to support the crucible when heated. Silicon dioxide is very 
inactive, but is readily attacked by hydrofluoric acid (p. 205). 
When heated to high temperatures with compounds of the 
metals, especially the carbonates and sulfates, silicates (salts 
of silicic acids) are formed as follows : 

Na 2 C0 3 -{- Si0 2 — ► Na 2 Si0 8 + C0 2 
Na 2 S0 4 + Si0 2 — > Na 2 Si0 8 + S0 3 

It is this reaction which is used in making glass. 




Fig. 170. A cluster of quartz crystals 






SILICON AND BORON 



289 



Silicic acids. Silicon forms two simple acids, orthosilicic acid 
(H 4 Si0 4 ) and metasilicic acid (H 2 Si0 3 ). Orthosilicic acid is 
formed as a jellylike mass when orthosilicates are treated with 
strong acids. If one attempts to dry this acid it loses water, 
passing into metasilicic acid 
(common silicic acid) : 

H 4 Si0 4 — >H 2 SiO s + H 2 

Metasilicic acid, when heated, 
breaks up into silica and water, 
thus : 



H 2 Si0 3 



H 2 + SiO, 



Salts of silicic acids ; sili- 
cates. A number of salts of 
the orthosilicic and metasilicic 
acids occur in nature. Thus, 
mica (KAlSi0 4 ) is a salt of 
orthosilicic acid. 




Fig. 171. A crucible and a triangle 
made from quartz 



Polysilicic acids. Silicon has 
the power to form a great many 
complex acids which may be regarded as derived from the union 
of several molecules of orthosilicic acid, with the loss of water. 
These are called polysilicic acids. For example, we have 



3 H SiO A 

4 4 



H 4 Si 3 8 + 4H 2 



Salts of these acids make up the great bulk of the earth's crust. 
Feldspar, for example, has the formula KAlSi 3 O g and is a salt of 
the acid H 4 Si 3 8 , whose formation is represented in the equation 
above. Kaolin, or pure clay, has the formula H 4 Al 2 Si 2 9 or, as 
commonly written, Al 2 Si 2 7 • 2 H 2 0. Granite is composed of crys- 
tals of feldspar and mica cemented together with amorphous silica. 

Water glass. Sodium silicate and potassium silicate differ 
from the silicates of the other metals in that they are soluble 
in water. The common water glass is a solution of sodium 
silicate. It is a thick, sticky liquid made by fusing sand with 



290 



CHEMISTRY AND ITS USES 



sodium carbonate. It is used for the purpose of giving a glazed 
waterproof surface to porous materials, such as wood, stone, 
and plaster ; to render curtains noninflammable ; as a glue 
for mending glass and pottery and for making strawboard 
boxes ; and as an ingredient of cheap soaps." 

Its property of closing pores is turned to account in preserving 
eggs for winter use. The eggs are packed in crocks and then 
covered with a liquid made by adding 1 volume of commercial 
water glass to 10 volumes of water. Over the liquid is then poured 
a little melted paraffin, which soon hardens and excludes the air. 
Fresh eggs can easily be preserved for six months in this way. 



General discussion 
pare, and in the free 



Boron 

Boron is a gray solid, difficult to pre- 
state but little used. It occurs in nature 
in the form of the mineral cohmanite 
(<Wy5H,0) (Fig. 172). This 
is mined in southern California and is 
the source of the compounds of boron 
prepared for commercial uses. 

Compounds. The most important 
compounds of boron are the following : 

Boric acid (H 3 B0 3 ). This is a white 
solid, slippery to the touch. It is a mild 
antiseptic and is used in medicines. 

Sodium tetraborate (Na 2 B 4 7 ). If 
we add sodium hydroxide to boric acid 

we do not get a salt of the simple 
Fig. 172. A specimen of acid? but one that has the f ormu l a 
colemanite, the mineral A -r-r»/-v i • n *» t . i 

from which borax is made Na 2 B 4° 7 and 1S called sodmm tetraborate. 

When this is crystallized from water at 
ordinary temperatures the crystals are composed of this salt, 
together with water, and have the formula Na 2 B 4 ? • 10 H 2 0. 
This compound is called borax. 







SILICON AND BOROX 



291 



Water of crystallization ; hydrates. Many salts when crys- 
tallized from water are found to have formed compounds 
with the water, and if the crystals are heated the water can 
be driven off as steam. Such salts are called hydrates, and 
the water is sometimes referred to as water of crystallization 
or water of hydration — though many 
crystals do not contain water. A salt -* 

which has crystallized without combin- 
ing with water is said to be anhydrous. 
In the formula of a hydrate it is cus- 
tomary to write the formula of the salt 
and the water separately with a period 
between. Thus the formula of borax is 
written Na 2 B 4 7 . 10 H 2 0. This formula 
expresses the fact that in the crystals 
each molecule of sodium tetraborate is 
combined with 10 molecules of water. 

Borax (Na 2 B 4 7 • 10 H 2 0). Borax (Fig. 
173) is the most important compound 
of boron, many thousands of tons of it 
being made annually from colemanite. 
It is a white solid and crystallizes in 
large crystals. When borax is heated, it 
swells up in a sort of froth, owing to 
the escape of steam, and this soon melts 
to a clear glass. The glass has the prop- 
erty of easily dissolving many metallic oxides, and this fact 
is turned to account in working with metals. When two 
pieces of metal are to be joined by melting them together or 
by the use of some kind of solder, the surfaces must be clean 
and free from oxide. Brass is joined by melting borax over 
the joint to clean the metal and then using a low-melting 
brass as a solder (brazing). Metallic oxides dissolved hi 
melted borax often color the borax glass with characteristic 




Fig. 173. Large crys- 
tals of borax 



292 CHEMISTRY AND ITS USES 

tints. On this account little beads of borax are used in test- 
ing for the presence of such metals. 

Borax is extensively used as a constituent of enamels and 
glazes for both metal ware and pottery. It is often used in 
our homes to soften hard water, as a mild alkali (like soap), 
and as an antiseptic. 

EXERCISES 

1. What substances have we studied that are prepared (a) by heat- 
ing in an electric furnace? (b) by passing a current of electricity 
through a liquid? 

2. How could you distinguish between a quartz crystal and a 
diamond? 

3. How do you account for the fact that some samples of sand are 
white and others colored ? 

4. What is the formula of the acid of which kaolin is the salt? 

5. What is the advantage and the disadvantage of quartz glass? 

6. To what class of substances does granite belong ? 

7. Is the water of hydration present in a compound in the form of 
water or is the water in a combined state ? 

8. Could you tell from the appearance of a compound whether it 
contained water of hydration? 

9. What is the weight of the water of crystallization in 1 ton 
of borax? 

10. State the chemical composition of each of the following sub- 
stances : graphite, boneblack, muriatic acid, caustic soda, kaolin, sand, 
flint, opal. 

11. Give another name for each of the following: laughing gas, oil 
of vitriol, cane sugar, cream of tartar, white arsenic. 






CHAPTER XXXII 



COLLOIDS - THE CHEMISTRY OF VERY SMALL PARTICLES 



Introduction. We all know that coarse sand shaken up 
with water quickly settles when we stop shaking the mix- 
ture. The finer the sand the slower it will settle. If we 
powder it fine enough and then stir it into water, it will not 
settle at all, for the particles will be so small that they will 
be kept in constant motion by collision with the molecules of 
the water. If we could powder the sand into individual mole- 
cules and stir them into water, we would have a true solution. 

In this chapter we shall be in- 
terested in these very fine particles 
that are too small to settl'e, cannot 
be seen directly by even the best 
of microscopes, and yet consist of 
many thousands of molecules. Such 
particles suspended in a liquid are 
called colloidal dispersions or col- 
loids^ and matter in this condition 
is said to be in the colloidal state. 

The air in a quiet room contains FiG m A beam of light shin . 
many particles of colloidal size that i ng through a colloidal solution 
are too small to be seen as objects by 

a microscope. But if the room is dark and a beam of sunshine 
enters through a small hole, these minute particles can be seen as 
bright, flashing points or motes dancing in the sunbeam. In like 
manner a colloidal suspension in a liquid that appears perfectly 
clear to the eye is seen to be full of moving particles when a 
strong beam of light shines through the liquid (Fig. 174), 

293 




294 



CHEMISTKY AND ITS USES 




How colloids are made. There are two general plans for 
making colloids, each of which may be modified in many 
details : 

1. An insoluble solid may be actually powdered to the 
necessary fineness and then stirred into the liquid (usually 
water). For example, we may prepare a colloidal dispersion 
of gold, silver, or platinum by establishing an electric arc 
between wires of the metal immersed in water. The current 

tears off minute particles as a 
kind of smoke and disperses them 
through the water (Fig. 175). 
Almost any sample of sand or 
ocean water shows traces of gold 
probably powdered to colloidal 
size by the natural grinding proc- 
esses of nature. 

2. We may bring about a re- 
action that ordinarily produces a 
precipitate, but under conditions 
such that the minute particles formed at first fail to gather 
together into particles large enough to settle as a precipitate. 
For example, we may add acid to a dilute solution of the salt 
known as sodium thiosulfate. No immediate precipitate of 
sulfur forms, but colloidal sulfur is present, as is evident 
from the white color. After a time some sulfur may precipi- 
tate, but most of it remains in suspension. 

Properties of colloids. Most of the interesting properties of 
colloids are due to the size of the particles rather than to the 
kind of matter of which they are made. 

1. All colloidal particles are charged electrically, some kinds 
positively and some negatively, and in an electrical field they 
act like the ions of a salt. 

2. Many colloids are highly colored, but the color gives us 
little clue as to what is present. Thus colloidal gold may 



Fig. 175. The preparation of 

metallic colloids by sparking 

under water 



COLLOIDS 295 

be red, blue, green, or violet, depending upon the size and 
uniformity of the particles. A little ferric chloride added to 
boiling water gives a deep-red colloidal iron oxide. It takes 
very little of the colloidal substance to make an intense color. 
A few milligrams of gold will color a liter of water. A very 
little of the element selenium added to glass makes the in- 
tensely red glass of the automobile tail-light. Many colored 
glasses and glazes owe their color to colloidal material of vari- 
ous sorts, and many gems are colored in the same way. The 
blue of the sky and of deep mountain lakes is probably due 
to colloidal dust in the air and the water. 

3. Since the colloidal particles are exceedingly small they 
cannot be filtered out by any ordinary filter. 

4. Many colloids have the property of adsorbing dissolved 
material from solutions. Thus ammonia is easily adsorbed by 
various colloids. 

Coagulation of colloids. It often happens that a colloidal 
dispersion is objectionable, and w T e want to remove it or re- 
cover the colloidal material. We must then change the elec- 
trical or chemical conditions so that the very small particles 
will grow together mto large clumps that will precipitate. 
We have found a number of ways to accomplish this : 

1. A given colloid is apt to remain suspended in an alka- 
line solution but to be precipitated in an acid solution ; or it may 
be just the reverse. So the addition of a little acid or alkali is 
apt to cause the colloid to precipitate, or coagulate, as it is called. 

Thus casein, an important colloid in milk, coagulates to a curd 
when the milk becomes sour or acid in reaction. Rubber is coagu- 
lated from the sap of rubber trees by the addition of acetic acid. 
If we try to prepare antimony sulfide in alkaline solution by bring- 
ing together hydrogen sulfide and a salt of antimony, we get only 
a deep-red colloidal dispersion, but if we add acid we get a brick- 
red precipitate. On the other hand, . many metallic oxides are 
precipitated by alkalies. 



296 CHEMISTRY AND ITS USES 

2. Sometimes a concentrated salt solution will coagulate 
a colloid. Thus most river water carries much colloidal clay 
and organic matter. When it reaches the salt water of the 
ocean this matter is precipitated and builds up the great river 
deltas such as those of the Mississippi or the Nile. In every 
soap factory colloidal soap is separated from its solutions by 
the addition of common salt. 

3. We have seen that all colloids are electrically charged. 
It has been found that when two colloids of opposite charge 
are brought together, they mutually precipitate each other. 
Thus the negative colloids of river water are precipitated by 
the positive colloids supplied by alum or iron salts in ordinary 
water purification. The processes of dyeing and tanning 
depend upon this same principle. 

Dispersal of colloids. It frequently happens that in many 
industrial operations we want to preserve the colloidal state 
or bring a colloidal precipitate once more into dispersed con- 
dition. For many purposes the colloidal suspension of an 
insoluble solid is just as good as a true solution. We can 
sometimes use the original methods for preparing a colloid. 
Thus, by a process similar to grinding, graphite may be sus- 
pended in oil or water to make a very effective lubricant for 
machinery (oil dag and water dag). We may reverse the 
acidity or alkalinity that caused the coagulation. Thus, anti- 
mony sulfide precipitated in acid solution may be dispersed 
again in an alkaline medium. We may filter the precipitate 
from the salt solution that produced precipitation and dis- 
perse it once more in pure water. 

If we bring together two colloids of the same electrical sign 
they tend each to hold the other in the colloidal state. Thus, 
we add gum arabic to some kinds of inks to hold the black 
ink material in suspension. This is the chief reason for adding 
gums, dextrin, starch, and similar materials to many industrial 
and medicinal preparations. 



COLLOIDS 



297 




Colloids and crystallization. It is an interesting fact that the 
presence of a colloid often tends to prevent distinct crystallization 
or at least to make the crystals very small and soft. Hence we put 
colloidal gelatin into ice cream, and gelatin, starch, dextrin, and 
similar materials in many candies and in electrolytic baths during 
electroplating, for in these processes 
we wish to avoid large or hard crys- 
tals. The colloidal materials in honey 
keep the concentrated sugar solution 
from crystallizing. 

Emulsions. If we pour two 
liquids together that do not mix, 
like oil and water, and shake the 
mixture violently, we get a milky- 
looking fluid called an emulsion. 
This consists of very minute drops 
of the one liquid dispersed through 
the other. If we let the emulsion 
of oil in water stand, it will very 
soon separate into the original 
liquids from which it was made. 
To make an emulsion of any kind 
permanent we must add a third 
colloidal substance insoluble in 
both of the liquids. This is called 
the emulsifying agent (Fig. 176). 
It seems to form a little skin over 
the surface of the drops and pre- 
vents them from running together into big ones. Milk is an 
emulsion of butter fat in water, with casein as the emulsi- 
fying agent. When the milk turns sour and the casein is 
coagulated, the butter fat is easily collected into large lumps 
of butter. Most of our common disinfectants and sheep dips 
owe their properties to cresol emulsified in water, with a small 
percentage of soap as the emulsifying agent. In mayonnaise, 




A B 

Fig. 176. Emulsions 

When oil and water are shaken 
together, the oil soon separates 
on standing (A) ; if a little soap 
is added to the oil and water, and 
the mixture shaken, the emulsion 
becomes more permanent (B) 



298 



CHEMISTRY AND ITS USES 



olive oil is emulsified in water (usually with a little vinegar 
added) by the colloidal yolk of egg, forming a stiff almost 
jellylike product. The fluids that form the webs of spiders 
and silkworms are secreted as emulsions that almost at once 
dry to form very fine filaments. 

Jellies. Sometimes a colloidal substance tends to form very 
thin films rather than particles or filaments, and these films 




Fig. 177. View of a factory showing the effect of the Cottrell method for 
abating dust and smoke 

may form in such a way as to inclose the liquid much as the 
walls of the honeycomb inclose the honey or the pores of a 
sponge hold water. When this happens the whole dispersion 
is likely to set to a more or less firm form called a jelly or 
a gel. Thus, if we add acid to a solution of sodium silicate 
the dilute dispersion of silicic acid soon sets to a firm jelly. 
Casein sets to a jelly when coagulated (Fig. 144). In fruit 
jellies it is a constituent of the unripe fruit called pectin that 
serves to form the supporting structure. In gelatin or glue 
the jelly may be dried out to a very compact form, but when 
dissolved in hot water and cooled a jelly is once more obtained. 
Soap is a partially dried jelly, and many minerals such as 



COLLOIDS 299 

agate, flint, and opal are dry jellies. Photographic films and 
many high explosives are jellies. Almost any sparingly soluble 
substance that does not crystallize too readily may be ob- 
tained in the form of a jelly. The so-called solidified alcohol, 
used largely at the present time, consists of a porous soap 
(Chapter XXXV) and alcohol, the soap acting as a sort of 
sponge to hold the alcohol. As yet we do not understand the 
conditions that lead to the formation of these jellies, and we 
obtain them largely by accident or by a cut-and-try process. 

Smokes. Evidently we may have very fine particles sus- 
pended in the air rather than in water, and such a suspension 
is a smoke. It is often very hard to condense the particles of 
a smoke into solid form, and this fact is of enormous industrial 
importance. Many precious materials literally " go up in 
smoke " from the stacks of smelters ; much zinc is lost as a 
smoke of zinc oxide in brass foundries ; valuable potassium 
compounds are lost in the smoke of cement burners ; and 
ordinary carbon smoke causes great losses to the individual, 
but profit to the laundries. 

All these smoke particles are electrically charged, and the 
American chemist Cottrell has devised a most effective process for 
the recovery of smoke particles by causing the smoke to move past 
metal plates and points charged electrically. The smoke particles 
lose their charge and then coagulate much as a precipitate and 
settle on the sides of the chimney (Fig. 177). The tails of comets 
are largely made up of particles of colloidal size. 

Fogs and foams. Fogs consist of very fine droplets of liquid 
(usually water) suspended in air, so that they correspond to emul- 
sions. Foams are gas bubbles dispersed through a liquid. Some- 
times these minute gas bubbles are dispersed through what we 
should usually consider a solid, and often give to it a pure-white 
color. Many white flowers, like lilies, owe their whiteness to dis- 
persed gas bubbles, and it appears that white hair is also due to air 
bubbles in the hair, though no one knows why age or worry should 
affect the hair in such a way. 



300 CHEMISTKY AND ITS USES 

EXERCISES 

1. Keeping the kinetic theory in mind (p. 47), explain why large 
particles suspended in a liquid will settle and very small ones will not. 

2. (a) What is the difference between saying that a solution is 
clear and that it is colorless? (b) Are colloid suspensions ever clear? 
(e) Are they colorless? 

3. The famous Hope diamond is very blue in color. Would a 
chemical analysis necessarily show the presence of a color pigment? 

4. The white of an egg is a typical colloid. In what way is it often 
coagulated? Can it be restored to the colloidal state? 

5. After oil is used in machinery it gets very black (as in the 
crank case of an automobile), yet very little of the color is removed by 
ordinary nitration or settling. How do you account for this fact? 

6. Why should a line suspension be as good as a true solution for 
many purposes? 

7. Many salves and ointments are made by rubbing an oil or a 
soft fat with some other material. How would you classify these 
preparations ? 

8. It is difficult to make a good jelly from ripe fruit. Can you 
suggest a reason for this? 

9. There is often enough sugar in a good fruit jelly to make a 
supersaturated solution. Why does it not more often crystallize out 
or ** sugar " ? 

10. Fertile soils are essentially a kind of colloid jelly, made up of 
clay and organic matter. When soluble sodium nitrate is used as a 
fertilizer why does not the first heavy rain wash it all away ? 

11. How do you account for the fact that silicon dioxide occurs in 
nature in so many different colors (p. 287) ? 

12. Ice made directly from natural waters is~ordinarily white, while 
that made from distilled water is clear and colorless. Suggest an 
explanation for this difference. 



CHAPTER XXXIII 
THE METALS 

Metals. The elements so far considered have nearly all been 
those whose compounds with oxygen and hydrogen are acids. 
They are called the acid-forming elements. Those which 
we shall now study are known collectively as the metals. 
Their hydroxides are bases, and on this account the metals may 
be defined as those elements whose hydroxides are bases. When 
the hydroxide of a metal or any of the simple salts are dis- 
solved in water, the metal forms the positively charged ion. 

Properties of the metals. The metals are all solids, except 
mercury, which is a liquid. Most metals have a high density, 
are good conductors of heat and electricity, are notably crystal- 
line in structure, and take a bright polish. With the exception 
of gold and copper they hav,e a silvery luster. Most of them 
combine readily with oxygen and sulfur, and their surfaces 
quickly tarnish on exposure to the air. 

A few of the least active of the metals, such as gold and 
copper, occur to some extent in nature in the native state. 
Most of them are found in combination as oxides, as sulfides, 
or as salts of various acids, especially as silicates. The process 
of winning metals from their ores is called metallurgy. The 
details of the metallurgy of any given metal must be adapted 
to the properties of the metal, to the form in which it is found, 
and to the cost of the operation. 

Compounds of the metals. Since the metallic elements are 
base-forming elements, the compounds which they form are 
chiefly oxides, hydroxides, and salts of various acids. We 

301 



302 CHEMISTRY AND ITS USES 

have seen (p. 146) that the number of salts is very large. It 
will be possible for us to study only the most important of 
these. With very few exceptions the compounds of the 
metals are all solids. 

Preparation of salts. There are two general ways of pre- 
paring salts which are employed so often that it is well to 
fix them in mind at the outset of the study of metals. 

1. Soluble salts. If a given salt is soluble it can usually be 
prepared in solution by treating the proper metal, or its oxide, 
hydroxide, or carbonate, with the proper acid. All these 
reactions have already been illustrated repeatedly : 

Cu + 2 H 2 S0 4 — y CuS0 4 + S0 2 + 2 H 2 (p. 188) 

CuO + H 2 S0 4 — >- CuS0 4 + H 2 (p. 193) 

NaOH + HC1 — > NaCl + H,0 (p. 145) 

CaC0 8 + 2 HC1 — >■ CaCl 2 + H 2 + C0 2 (p. 112) 

2. Insoluble salts. A very large number of salts are insoluble 
in water, and these can be made by precipitation. All that it 
is necessary to do is to bring together in solution two salts, 
one of which contains the desired metallic ion and the other 
the negatively charged ion.. If, for example, it is known that 
silver chloride is insoluble in water, it is to be expected that 
this salt will be precipitated on mixing a solution of any silver 
salt, such as silver nitrate, with that of any chloride, such as 
sodium chloride, thus : 

AgN0 8 + NaCl — >- AgCl + NaN0 3 

Electrochemical industries. To an ever-increasing extent 
electrical energy is being used both in the separation and the 
refining of the metals and for the production of. many com- 
pounds. Such methods have in a number of instances already 
been mentioned. Naturally these industries tend to develop 
most extensively in localities where water power is abundant. 



THE METALS 303 

Norway has many electrochemical industries. Those of the 
United States and Canada center at Niagara Falls, the exten- 
sive power plants being shown in Fig. 178. 

Solubility of salts. It will be seen that a knowledge of the 
solubility of various salts is of great importance if one wishes 
to devise a means of preparing a given salt. Fortunately it 




Fig. 178. Electrochemical power plants at Niagara Falls 

is possible to put into brief form the facts relating to the 
solubility of the common salts, and these rules will have 
constant application in the pages which follow. 

1. Hydroxides. All hydroxides are insoluble except those of 
ammonium, sodium, potassium, calcium, barium, and strontium. 

2. Nitrates. All nitrates are soluble. 

3. Chlorides. All chlorides are soluble except silver and mer- 
curous chlorides. (Lead chloride is soluble in hot water.) 

4. Sulfates. All sidfates are soluble except those of lead, 
barium, and strontium. (Sulfates of silver and calcium are 
only moderately soluble.) 

5. Sulfides. All sulfides are insoluble except those of am- 
monium, sodium, and potassium. The sulfides of calcium, 
barium, strontium, and magnesium are insoluble in water, 



304 CHEMISTRY AND ITS USES 

but are changed by water into acid sulfides which are solu- 
ble. On this account they cannot be prepared by precipitation. 
6. Carbonates, phosphates, and silicates. All normal carbonates, 
phosphates, and silicates are insoluble except those of ammonium, 
sodium, and potassium. 

EXERCISES 

1. Make a list of the elements so far studied. How many are gases 
under ordinary conditions ? How many are liquids ? 

2. What elements studied are prepared commercially through the 
agency of electricity ? 

3. Why cannot sodium be produced by the electrolysis of a solution 
of its salts ? In general, what metals could be so produced ? 

4. What is the connection between water power and the production 
of metals and their compounds? 

5. Suppose you have on hand hydrochloric acid and marble (p. 112) 
and wish to prepare some calcium chloride for drying gases (p. 35). 
How would you proceed ? 

6. Ordinary copperas has the formula FeS0 4 • 7 H 2 0. To what class 
of substances does this compound belong ? What materials would you 
require for its preparation ? 

7. Suppose you wished to prepare some potassium iodide for medici- 
nal use and had on hand potassium hydroxide, (a) What other compound 
would you require ? (A) Describe the process, (c) Show that the reaction 
you employ would go to completion (or nearly so). 

8. Ferric hydroxide (Fe(OH) 3 ) is a reddish solid, (a) Is it soluble 
in water (p. 303)? (b) How could you convert the copperas, prepared 
accordiug to Ex. 6, into ferric hydroxide ? 

9. Suppose you mixed solutions of sodium carbonate (Na 2 C0 3 ) and 
calcium chloride (CaCl 2 ). (a) Should you expect any reaction to take 
place? (b) Explain. 

10. The hydrate of copper sulfate (CuS0 4 • 5 H 2 0) is commonly known 
as blue vitriol Suggest a method for making it. 



CHAPTER XXXIV 
THE SODIUM FAMILY 



Name 


Symbol 


Atomic _ 

WEIGHT ) DENSITY 


Melting 
Point 


First prepared 


Lithium . . 
Sodium . . 
Potassium . . 
Rubidium . . 
Caesium . . 


Li 

Na 

K 

Rb 

Cs 


6.94 

23.00 

39.10 

85.45 

132.81 


0.534 
0.971 
0.862 
1.53 

1.87 


186° 

97° 
62° 
38° 
26° 


Arf vedson, 1817 
Davy, 1807 
Davy, 1807 
Bunsen, 1861 
Bunsen, 1860 



The family. These elements, known as the very light 
metals, constitute a family in Group I in the periodic table. 
The name alkali metals is often applied to the family for the 
reason that the hydroxides of the most familiar members of 
the family (namely, sodium and potassium) have long been 
called alkalies. Sodium and potassium are the only important 
members of the family. We have already studied sodium 
(Chapter XVI), and it is advisable for the student to review 
that chapter in connection with the present one. 

Compounds of Sodium 

Introduction. Everyone is familiar with some of the metals 
such as silver, lead, copper, and iron because of their use as 
metals in our daily lives. In other cases few of us have 
ever seen the metal, but its compounds are very well known. 
Sodium is a metal of this kind. As a metal it is a curiosity 
to most of us, but we all know something about salt, lye, soda, 
soap, and glass, and these are all compounds of sodium. We 
have already studied sodium hydroxide as an example of a 

305 



306 CHEMISTRY AND ITS USES 

typical base (Chapter XVI). We shall now study some other 
important compounds of sodium. Nearly all these are white 
crystalline solids. 

Sodium chloride (common salt) (NaCl). Sodium chloride, or 
common salt, is very widely distributed in nature. Thick 
strata, evidently deposited by the evaporation of salt water, 
are found in many places. In the United States the most 
important localities for salt are New York, Michigan, Ohio, 















f': > r; 


fyX 










l 


mm 




»r mm m% 








89 — ; . . 








■:■■■■.■■■ ■. 

1 







Fig. 179. The evaporation of salt brine in the open air 

and Kansas. Sometimes the salt is mined (Fig. 81), espe- 
cially if it is in the pure form called rock salt. More fre- 
quently a strong brine is pumped from deep wells sunk into 
the salt deposit. The brine is evaporated either by heating 
or, in the preparation of the coarser grades of salt, by simply 
exposing the brine to the air (Fig. 179). Salt crystallizes in 
the form of small cubes. 

Uses of salt. Since salt is so abundant in nature it is used 
in the preparation of nearly all compounds containing either 
sodium or chlorine. These include many substances of the 
highest importance to civilization, such as soap, glass, hydro- 
chloric acid, soda, and bleaching powder. Enormous quantities 
of salt are therefore produced each year. Small quantities are 
essential to the life of man and animals. 



THE SODIUM FAMILY 307 

Pure salt does not absorb moisture ; the fact that ordinary salt 
becomes moist when exposed to air is not due to a property of the 
salt but to impurities occurring in it, especially to the presence of 
the chlorides of calcium and magnesium. 

Sodium sulfate (Na 2 S0 4 ). This salt is prepared by the action 
of sulfuric acid upon sodium chloride, hydrogen chloride 
being formed at the same time (p. 143) : 

2 NaCl + H 2 S0 4 >- Naj30 4 + 2 HC1 

The anhydrous salt is a white solid. It is readily soluble in 
water, and under ordinary conditions crystallizes as the 
h}'drate Na 2 S0 4 • 10 H 2 (known as Glauber s salty. Large 
quantities of sodium' sulfate are used in making sodium car- 
bonate and glass. 

Sodium carbonate (soda ash) (Na 2 C0 3 ). This is the sodium 
salt of the very weak, unstable carbonic acid (p. 113), but 
while the acid is unstable, the salt is very stable and is of 
great importance. There are two different methods now em- 
ployed in its manufacture, each of which bears the name of 
its inventor (Figs. 180, 181). 

1. Leblanc process. This older process, still used in England, 
involves several distinct reactions, as shown in the following 
equations : 

a. Sodium chloride is first converted into sodium sulfate : 

2 NaCl + H 2 S0 4 > Na 2 S0 4 + 2 HC1 

5. The sodium sulfate is next reduced to sulfide by heating 
it with carbon : 

Na 2 S0 4 -f 2 C — y Na 2 S -f 2 C0 2 

c. The sodium sulfide is then heated with calcium carbonate, 
when the following reaction takes place : 

Na 9 S + CaCO, ►■ CaS + Na 2 C0, 



308 CHEMISTRY AND ITS USES 

2. Solvay process. This newer process, and the only one used 
in the United States, consists in passing carbon dioxide and 
ammonia into a saturated solution of sodium chloride : 

a. NH 3 + C0 2 + NaCl + H 2 y NaHC0 3 + NH 4 C1 

The sodium hydrogen carbonate is then filtered off and heated: 

b. 2 NaHC0 3 > Na 2 C0 3 + H 2 + C0 2 

The ammonium chloride formed in equation a is utilized in 
once more preparing ammonia. 

When sodium carbonate is crystallized from water it forms 
large crystals of the formula Na 2 C0 3 • 10 H 2 0, usually known 
as washing soda or sal sotfa. Its solution in water has a mild 
basic reaction and is used for laundry purposes. Mere mention 
of the fact that it is used in the manufacture of glass, soap, 
and many chemical reagents will indicate its importance in the 
industries. It is one of the few soluble carbonates. 

Historical. In former times sodium carbonate was made by 
burning seaweeds and extracting the carbonate from their ash. 
On this account the salt was called soda ash, a name still in 
common use. During the French Revolution this supply was 
cut off, and in behalf of the French government Leblanc (Fig. 180) 
made a study of preparing the carbonate directly from salt. As a 
result he devised the method which bears his name and which 
was used exclusively for many years. It has been replaced to a 
large extent by the Solvay process, which was devised by the 
Belgian chemist Solvay (Fig. 181) in 1863. 

By-products. The substances obtained in a given process, aside 
from the main product, are called the by-products. The success of 
many processes depends upon the value of the by-products formed. 
Thus hydrochloric acid, a by-product in the Leblanc process, is 
valuable enough to make the process pay, even though sodium 
carbonate can be made more cheaply in other ways. 

Hydrolysis of salts. In connection with neutralization 
(p. 154) we learned that when an acid and a base are brought 




Fig. 180. Nicolas Leblaxc (1742-1806) 

This cut is from a statue erected in Paris in honor of Lehlanc. 
Previous to his time the supply of sodium carbonate came en- 
tirely from the ash of seaweeds. Lehlanc devised a method for 
making it from common salt. This proved to be a much cheaper 
source of the compound. The Lehlanc method was long used 
in the United States and is still used in England 







Fig. 181. Ernest Solvay 




■HBHHH 




HSHf§§Vr 


(1835- ) 
A famous Belgian manufactur- 




j|J| 


ing chemist who developed the 
Solvay process for making so- 




■ 


dium carbonate. The commercial 




H 


success of this process brought 






Solvay great wealth, and he has 




ii$? -* 


been very generous, distributing 




A • Jh^ 


on his seventy-fifth birthday 




jk ■■•"■'■H^fc_ 


over a million dollars for educa- 




Hi 


tional and philanthropic purposes 

■ 




• UkI/m^^II 


• 

Fig. 182. Victor Grignard 
(1871- ) 




't t* wHm* *** 






iw^PB^y^i 


A famous French chemist. He 




discovered a method for making 




iflwB 


a large number of useful organic 
compounds, for which discovery 
he was awarded the Nobel prize 
in 1912. During the World War 






the French government sent him 






to the United States to confer 




with our chemists on matters 




dB 


pertaining to the war 







THE SODIUM FAMILY 309 

together in solution the hydroxyl ion of the base and the 
hydrogen ion of the acid unite to form water, leaving the ions 
of the salt, which form a neutral solution. From this it might 
be inferred that solutions of all salts in water are neutral to 
indicators. This is true if both add and base from which the 
salt is formed are strong ones. If either is very weak, then the 
water acts upon the salt and to a slight extent reverses the 
action of neutralization. Thus in the case of sodium carbon- 
ate we have a slight reaction as follows : 

Na 2 C0 3 + H 2 :<=>: NaOH + NaHC0 3 

The action of water on a salt to form an arid and a base is called 
liydrolysis. Since the hydrogen of the NaHC0 3 is a part of 
the weak acid H o C0 8 , it is little ionized, while the NaOH 
formed is a strong base and is very largely ionized. Conse- 
quently the resulting solution is alkaline in reaction. In general, 
all sodium and j^otassium salts of weak acids are alkaline in 
reaction. For a similar reason all salts of a weak base with 
a strong acid are acid in reaction. 

Sodium hydrogen carbonate (NaHC0 3 ). This salt, called 
bicarbonate of soda or baking soda, is made by the Solvay proc- 
ess, as explained above, or by passing carbon dioxide into 
concentrated solutions of sodium carbonate : 

Na,CO, + H o + CO, y 2 NaHCO, 

It is an essential constituent of all baking powders. 

Sodium nitrate (NaN0 3 ). This substance, known also as 
Chile saltpeter, is found in nature in certain arid regions, 
where apparently it has been formed by the decay of organic 
substances in the presence of air and sodium salts. The larg- 
est deposits are hi Chile (Fig. 183), and most of the nitrate 
of commerce comes from that country. Smaller deposits occur 
in California and Nevada. The commercial salt is prepared 



310 



CHEMISTRY AND ITS USES 



by dissolving the crude nitrate (known as caliche) in water, 
allowing the insoluble earthy materials to settle, and evaporat- 
ing to crystallization the clear solution so obtained. The solu- 
ble impurities remain for the most part in the mother liquors. 
Since this salt is the only nitrate found extensively in nature, 
it is the material from which other nitrates, as well as nitric acid, 
are prepared. Enormous quantities are used as a fertilizer. 




Fig. 183. Deposit of sodium nitrate in Chile 



Other compounds of sodium. Among the other important com- 
pounds of sodium the following may be mentioned : (1) Sodium 
thiosulfate (Na 2 S 2 3 • 5 H 2 0), commonly known as sodium hypo- 
sulfite or, in photography, simply hypo. It is used in photography 
to dissolve off the plate or film those silver salts that have not 
been acted upon by the light. (2) Sodium cyanide (ISTaCN) is used 
in preparing prussic acid (p. 174). Since it readily dissolves gold, 
it is used in the extraction of gold from its ores. (3) Disodium 
phosphate (Na 2 HP0 4 ) is the most common of soluble phosphates. 
(4) Sodium peroxide (Na 2 2 ) is a strong oxidizing agent and 
evolves oxygen when brought in contact with water. (5) Sodium 
iodide (Nal) and sodium bromide (NaBr) are used in medicine and 
photography. (6) Sodium sulfite (Na 2 S0 3 ) is used as a preservative. 
Thus, when added to ground meat, it not only preserves the meat 
but causes it to retain the red appearance of fresh meat. 



THE SODIUM FAMILY 311 

Potassium 

Preparation and properties. Potassium is prepared by 
methods similar to those used in the preparation of sodium. 
It is more active than sodium ; otherwise the properties of 
the two metals are nearly alike. 

Occurrence. Potassium is a rather abundant element, being 
a constituent of many igneous rocks, such as the feldspars 
and micas. Very large deposits of the chloride and the sulfate, 
associated with compounds of calcium and magnesium, are 
found at Stassfurt, Germany, and are known as the Stassfurt 
salts. Similar deposits occur in Alsace, France. 

The natural decomposition of rocks containing potassium 
gives rise to various compounds of the element in all fertile 
soils. Its soluble compounds are absorbed by growing plants 
and built up into complex vegetable tissues ; when these are 
burned, the potassium remains in the ash in the form of car- 
bonate. The crude carbonate can be separated from wood 
ashes by dissolving it in water. This was formerly the chief 
source of potassium compounds, but they are now prepared 
mostly from the natural deposits in Germany and France. 

The natural deposits consist of definite minerals (Fig. 184), 
the most important of which are the following : 

Sylvite . . KC1 

Carnallite KC1 • MgCl 2 • 6 H 2 

Kainite KC1 • MgS0 4 • 3 H 2 

Kieserite MgS0 4 • H 2 

Compounds. The compounds of potassium are so similar in 
properties to the corresponding compounds of sodium that 
for many purposes for which they are used they can be inter- 
changed. The compounds of potassium, being as a rule more 
expensive, are not so widely used as those of sodium. The 
methods of preparation are in general the same as those used 
in preparing the corresponding compounds of sodium. 



312 



CHEMISTRY AND ITS USES 



Potassium chlorate (KC10 3 ). This salt can be made by the 
action of chlorine on hot solutions ol* potassium hydroxide. The 
reaction is somewhat complex, as shown by the following equa- 
tion 



3 Cl 2 + 6 KOH 



+■ 5 KC1 + KCIO, 



+ 3 H 2 





t 
r ~\ i 


^MMMUIr-v 


■H W 

HI 

1 KAlNiTE 1 


[.4 


^^^ 




1 SYLVITE M 


IcARNALUTEH 


U KIESER1TEH 


MsylvihiteH 




K — «4 


W^;*r^>-,;i 










i " i 


:: 










Sm-^Br 


w-JSlJ 


fc.v 




1 






•■:,:'' "" Si 


v:-c- — -2|' 



Fig. 184. The most important of the minerals constituting the 
Stassfurt salts 



It is a white crystalline substance and is used chiefly as an 
oxidizing agent in the manufacture of matches, fireworks, and 
explosives ; it is also used in the preparation of oxygen and 
in medicine. 

Potassium nitrate (saltpeter) (KN0 3 ). This salt is found 
in some regions where the climate is hot and dry, being formed 
by the decay of nitrogenous organic matter in the presence of 
earthy material containing compounds of potassium. The salt- 
peter used in making gunpowder was formerly made by imi- 
tating these conditions. At present it is prepared by the action 
of sodium nitrate upon potassium chloride (the former com- 
pound being obtained from Chile and the latter from Germany 
or France). 



NaNO, + KC1 



KNO. + NaCl 



THE SODIUM FAMILY 313 

This reaction can be made to work because, of the four com- 
pounds involved in the reaction, sodium chloride is the least solu- 
ble in boiling water and because its solubility is about the same 
in cold as in hot water (Appendix). Hence, if boiling saturated 
solutions of sodium nitrate and potassium chloride are mixed, 
sodium chloride will form and largely precipitate. When the 
solution is cooled, the potassium nitrate will then crystallize. 

Potassium nitrate is a white solid. It is an excellent oxidiz- 
ing agent and is chiefly used in the manufacture of gun- 
powder. Smaller amounts are used in medicine and as a 
preservative for meats. 

The question might naturally arise as to the reason for using 
potassium chlorate and potassium nitrate in place of the cor- 
responding compounds of sodium. The reason is that sodium chlo- 
rate is difficult to prepare pure because it is very soluble, while 
sodium nitrate attracts moisture from the air and therefore is not 
adapted for the manufacture of gunpowder, which must be kept dry. 

Other compounds of potassium. Potassium hydroxide or caustic 
potash (KOH) is a strong base. Potassium chloride (KC1), potas- 
sium bromide (KBr), and potassium iodide (KI) are all white solids. 
The bromide and the iodide are both used in the preparation of 
photographic reagents and in medicine. Potassium carbonate 
(K 2 C0 3 ), potassium bicarbonate (KHC0 3 ), potassium sulfate 
(K 2 S0 4 ), and potassium bisulfate (KHS0 4 ) are all well-known 
compounds. They are white solids, readily soluble in water. 

Sources of potassium salts in the United States. The great 
use of potassium salts is for fertilizer. Previous to the World 
War practically all our supply came from Germany. Indeed, 
when war was declared Germany counted on starving us, since 
we could not get these compounds from her. During the period 
of the war an effort was made to develop adequate supplies 
of potassium compounds within our borders. In this we were 
only partially successful, a certain amount being obtained 
from some of the salt lakes in Nebraska, from minerals, and 
from other minor sources. 



314 CHEMISTRY AND ITS USES 

Deliquescent compounds. We have seen that certain com- 
pounds such as calcium chloride (p. 35) and sodium nitrate 
attract moisture and become damp on exposure to moist air. 
Such compounds are said to be deliquescent. 

The Ammonium Compounds 

Composition. As explained in a previous chapter, when 
ammonia is passed into water the two combine to form the 
base NH 4 OH, known as ammonium hydroxide. When this base 
is neutralized with acids, there are formed the corresponding 
salts, known as the ammonium salts. Since the ammonium 
radical is univalent, ammonium salts resemble those of the 
alkali metals in their formulas ; they also resemble the latter 
salts very much in many chemical properties and may be 
conveniently described in connection with them. Among the 
ammonium salts the chloride, sulfate, carbonate, and sulfide 
are the most familiar. 

Ammonium chloride (sal ammoniac) (NH 4 C1). This is a white 
solid. When heated it partly decomposes into ammonia and 
hydrogen chloride, which recombine as the temperature falls : 

NH 4 C1^=>:NH 3 + HC1 

This salt is used in soldering, since the hydrogen chloride 
evolved by the heat removes the oxide from the surface of the 
metals. It is also used in making dry cells, in medicine, and 
as a chemical reagent. 

Ammonium sulfate ((NH 4 ) 2 S0 4 ). This salt is prepared at 
low cost by passing ammonia into sulfuric acid. It is espe- 
cially valuable as a fertilizer, and large quantities are used 
for this purpose. 

Carbonates of ammonium. Both the normal carbonate 
(NH 4 ) 2 C0 3 and the acid carbonate NH 4 HC0 3 are white solids, 
readily soluble in water. They are important chemical reagents. 



THE SODIUM FAMILY 



315 



Ammonium sulfides. When hydrogen sulfide is passed into 
aqua ammonia a solution containing ammonium sulfide 
((NH 4 ) 2 S) and ammonium acid sulfide (NH 4 HS) is obtained : 

NH OH + HS >■ NHHS + HO 



2 NH 4 OH + H 2 S 



> (NH 4 ) 2 S + 2 H 2 




This solution is usually known simply as ammonium sul- 
fide and is used in chemical analysis as a reagent in testing for 
certain metals. 

Flame reactions. A number of the metals when volatilized 
in a colorless flame, such as that of a Bunsen burner, impart 
a characteristic color to the flame. Thus sodium (or any of its 
compounds that will volatilize in the heat of the flame) makes 
the flame yellow. Potas- 



sium and its compounds 
color the flame a pale 
violet, and lithium colors 
it a deep crimson red. 

Advantage is taken of 
these facts in testing for 
the presence of the ele- 
ments in different sub- 
stances. The test is best 
made by using a platinum 
wire, one end of which is 

fused into a piece of glass tubing that serves as a handle. 
The other end of the wire is dipped into water and rubbed 
in the substance to be tested (or dipped into a concentrated 
solution of the substance), and the wire with the adhering 
particles is held in the outer edge of the base of the Bunsen 
flame (Fig. 185). The light from the sodium flame is absorbed 
in passing through blue glass, while that from potassium 
is not absorbed. Advantage is taken of this difference in 
detecting the two elements in mixtures. 



Fig. 185. Making a flame test 



316 CHEMISTRY AND ITS USES 

EXERCISES 

1. What is an alkali? Can a metal itself be an alkali? 

2. Do the alkali metals show any gradation of properties corre- 
sponding with their atomic weights ? (Consult table at head of chapter.) 

3. («) Which of the alkali metals will float on water ? (b) Which 
will melt in boiling water? 

4. How should you expect potassium to act on water? 

5. What substances should you expect to result from the electrolysis 
of a solution of potassium chloride ? 

6. What nonmetallic element is obtained commercially from the 
deposits of Chile saltpeter? 

7. How can you change sodium carbonate into sodium bicarbonate 
and vice versa? 

8. The sodium chloride used for table salt is often mixed with a 
little flour or starch. Explain the reason for this. 

9. Ammonium chloride is a by-product of the Solvay process. How 
could you prepare from this the ammonia used in the process (p. 164) ? 

10. Carbonic acid (H 2 C0 3 ) (p. 113) is a very weak and very unstable 
acid. How do you account for the fact that a solution of sodium car- 
bonate is basic? 

11. (a) Why are the carbonates such as sodium carbonate and cal- 
cium carbonate so easily acted upon by acids ? (b) How could you test 
for the presence of a carbonate? 

12. Write the equation for the reaction between sodium carbonate 
and hydrochloric acid. 

13. Why is sodium hydrogen carbonate used in fire extinguishers 
(p. 113)? 

14. The solution obtained by treating wood ashes with water is 
strongly basic. Explain. 

15. (a) What part did Chile saltpeter play in the World War? 
(b) How did Germany get along without it? 

16. (a) Suppose you passed chlorine into a hot solution of potassium 
hydroxide, what products are formed (p. 312) ? (b) If you were to heat 
the products, what single compound would be left ? 

17. Bromine and iodine act upon potassium hydroxide (and sodium 
hydroxide) just as chlorine does. How could you prepare («) potassium 
bromide ? (b) sodium iodide ? 



THE SODIUM FAMILY 317 

18. State what substances already studied are prepared from the 
following compounds: («) ammonium chloride; (b) ammonium nitrate : 
(c) sodium nitrate ; (d) potassium chlorate. 

19. 100 kg. of Glauber's salt contains what weight of water of 
crystallization ? 

20. 100 lb. of sodium carbonate is dissolved in water and crys- 
tallized, (a) What is the common name of the resulting product? 
(b) Calculate the weight of the product. 

21. If you used large quantities of sal soda and had to have it 
shipped a long distance, what economy would suggest itself to you ? 

22. Borax (p. 291) has a mild alkaline reaction. Can you explain 
this ? 

23. What weight of sodium carbonate can be prepared from 1 ton of 
sodium chloride («) by the Leblanc process (note that 2 gram-molec- 
ular weights of sodium chloride give 1 gram-molecular weight of sodium 
carbonate ; see equation, p. 307) ? (6) by the Solvay process ? 

24. Supposing that the fire extinguisher shown in Fig. 68 (p. 113) 
contained 5 kg. of sodium bicarbonate in solution, what volume of 
carbon dioxide would be evolved in using the extinguisher ? Equation : 

2 XaHC0 3 + H 2 S0 4 >- Xa 2 S0 4 + 2 H 2 + 2 C0 2 

25. Suppose you wished to operate a plant built for the manufacture 
of 3 tons of saltpeter daily, (a) What compounds should you require 
for its manufacture? (ft) Calculate the daily consumption of each of 
these compounds in your factory, assuming that you could obtain 
100 per cent yields. 

26. Suppose you were given 1000 g. of sal soda and told to con- 
vert it into caustic soda, (a) how would you proceed (p. 153) ? 
(b) what other materials would you require? (r?) what weight of 
sodium hydroxide would you expect to obtain ? 



CHAPTER XXXV 
SOAP; GLYCERIN; EXPLOSIVES 

Introductory. At first thought one might wonder why three 
products so different from each other as are soap, glycerin, 
and explosives should be brought together for study in one 
chapter. However, the grouping is a natural one industrially, 
for glycerin is a by-product in the manufacture of soap ; and 
nitroglycerin, one of the most powerful explosives, is prepared 
from glycerin and stands in a general way as a type of an 
explosive compound. 

Composition of soap, and materials used in its preparation. 
Ordinary soap is a mixture of the sodium salts of oleic, pal- 
mitic, and stearic acids (pp. 259, 262). The essential materials 
used in the preparation of soap are as follows : 

1. Fat or oil. As shown on page 262, fats and oils are 
largely mixtures of olein, palmitin, and stearin. The cheaper 
grades of these are used in making soap. Those commonly 
employed are a low grade of animal fat (tallow) and the 
cheaper vegetable oils, such as cottonseed oil, coconut oil, 
and palm oil (Fig. 186). Some resin (p. 408) is used in the 
cheaper grades of soap as a substitute for the fat. 

2. Alkali. The alkali used is the hydroxide of either sodium 
or potassium. Sodium hydroxide is nearly always used, since it 
gives a hard soap, while potassium hydroxide gives a soft soap. 

Reaction taking place in the preparation of soap. When the 
fat and alkali are heated together under proper conditions, 
the olein, palmitin, and stearin present in the fat are decom- 
posed, forming glycerin, together with sodium oleate, sodium 

318 



SOAP; GLYCERIN; EXPLOSIVES 



319 



palmitate, and sodium stearate. A mixture of these three salts 
constitutes soap. The reactions may be illustrated by the fol- 
lowing equation, which represents the change taking place 
when stearin is heated with sodium hydroxide : 

C 8 H 5 (C 18 H 35 2 ) 3 + 3 NaOH —*■ C 8 H 5 (OH) 3 + 3 NaOJB.0, 



5V 18 
stearin 



sodium stearate 



sodium hydroxide glycerin 

In this reaction the fat (stearin, palmitin, olein) is said 
to be saponified, and the process is known as saponification. 



^<_*sar 




Fig. 186. Some common substances used in the manufacture of soap 

Commercial manufacture of soap. The oil or melted fat is poured 
into large iron kettles together with a solution of sodium hydrox- 
ide containing about one fourth of the amount of alkali necessary 
to saponify the fat. As a rule the kettles are very large (Fig. 187), 
500,000 lb. or more of soap being made in some of them in a single 
heating. They are provided with coils of steam pipe for heating 
the mixture. The fat and alkali are stirred by forcing air or live 
steam into the bottom of the mixture. As the heating continues 
the remainder of the alkali is added. The reaction is complete in 
from two to five days. 

The soap is next removed, or salted out (p. 296), from the mix- 
ture. This process consists in adding salt and again heating. 
After a time the soap rises to the top of the liquid (or spent lye, 
as it is called). The soap so obtained is purified and then run into 



320 



CHEMISTRY AND ITS USES 



a mixing machine (crutcher). Here it is mixed with any appro- 
priate material which it is desired to add, such as perfume, borax, 
sodium silicate, or sodium carbonate. It is then run into large 
molds to harden, after which it is cut and pressed into cakes of 
the desired size. The glycerin formed in the reaction is separated 

from the spent lye by distillation. 

Varieties of soap. Transparent soaps 
are ordinarily made by dissolving 
soap in alcohol. The solution is fil- 
tered and the excess of alcohol re- 
moved by distillation. Castile soaps 
are made from mixtures of olive oil 
with cheaper oils. The color of mot- 
tled soaps is produced by the addition 
of ferrous sulfate, Prussian blue, or 
some similar pigment. Floating soaps 
owe their lightness to bubbles of air. 
Naphtha soaps contain from 5 per cent 
to 10 per cent of petroleum naphtha. 
Scouriny-soapjs contain from 5 per cent 
to 10 per cent of soap and from 80 
per cent to 90 per cent of some abra- 
sive material such as fine sand or 
volcanic ash. Sometimes a small per- 
centage of sodium carbonate is also 
present. Soap powders are, as a rule, sodium carbonate or a 
mixture of sodium carbonate and ground soap. 




Fig. 187. A large kettle for 
making soap 



Properties of soap. Soap forms with soft water a colloidal 
solution (p. 293). The resulting liquid is slightly alkaline 
due to hydrolysis. If an acid, such as hydrochloric acid, is 
added to the colloidal solution, the organic acids are liberated 
from their salts and, being insoluble, are precipitated : 



NaC 18 H 35 2 + HCl 

sodium stearate 



NaCl + H.C„H,,0, 

stearic acid 



The calcium and magnesium salts of oleic, palmitic, and stearic 
acids are insoluble in water and are therefore precipitated when 



SOAP; GLYCERIN; EXPLOSIVES 



321 



a calcium or magnesium compound is added to a colloidal 
solution of soap (Fig. 188) : 



2NaC 18 H 86 2 +CaCl 2 

sodium stearate 



2NaCl + Ca(C H 2 X 



calcium stearate 



It is due to this fact that soaps do not lather with hard waters 
but form a curdy precipitate, since such waters always contain 
calcium and mag- 
nesium salts in so- 
lution. 

The cleansing ac- 
tion of soap is not 
thoroughly under- 
stood, but seems to 
be associated with 
the property which 
soap possesses of 
forming emulsions 
(p. 297). When 
soap is rubbed on 
the skin any oily 
substances present 
are emulsified and 
washed away. 

Candles. When 
fats are heated with 
steam they are decomposed into glycerin and free acid as follows : 




Fig. 188. When solutions of soap and a calcium 

salt are mixed, a white precipitate of calcium 

stearate forms 



C s H 5 (C 18 H 36 2 ) 3 +3HOH 



C,H,(OH),+ 8H.C„H i ,0 1 



glycerin 



stearic acid 



The solid fatty acids thus obtained, mixed with paraffin, are 
used in making candles. The glycerin set free in this process, 
as well as that set free in the process of soap-making, is easily 
recovered, and it is in this way that our commercial supplies 
of glycerin are obtained. 



322 CHEMISTRY AND ITS USES 

Glycerin (C 3 H 5 (OH) 3 ). This is a colorless oily liquid having 
a sweet taste. Nitric acid reacts with it just as with a base, 
forming the nitrate C 3 H 5 (N0 3 ) 3 : 

C s H 5 (OH) s + 3 HN0 3 — y C 3 H 5 (N0 3 ) 3 + 3 HOH (H,0) 

The nitric acid used in this reaction is mixed with a little 
sulfuric acid, the latter serving to absorb the water produced. 
The resulting nitrate is a slightly yellowish oil commonly 
known as nitroglycerin. It is very explosive and is used in 
the manufacture of dynamite. 

Utilization of garbage. Many cities now treat their garbage in 
such a way that it becomes a source of revenue rather than expense. 
The garbage is first heated with steam and hot water. This 
causes the oils and fats to collect as a liquid on the top of the 
hot water so that they can easily be removed. The residue 
makes a valuable fertilizer, while the fats and oils are used in 
making candles, soap, and glycerin. 

Explosives. An explosion is caused by a very rapid chem- 
ical reaction which results in the formation of a large volume 
of gas from a comparatively small volume of liquid and solid 
materials called explosives. The greater the volume change and 
the more rapidly it is produced, the more violent the explosion. 

For a brief discussion we may divide explosives into two 
groups : namely, (1) gunpowder and (2) nitro-explosives. 

1. Gunpowder. Ordinary gunpowder is an intimate mixture of. 
potassium nitrate, sulfur, and charcoal. When they are ignited, 
complicated reactions occur and a number of products are formed, 
about half of which are solid and the remaining half gases. 
The chief of these products are carbon dioxide, nitrogen, and the 
carbonate, sulfate, and sulfide of potassium. 

2. Nitro-explosives. These are compounds of carbon, hydrogen, 
oxygen, and nitrogen. When they are exploded, the carbon and 
hydrogen unite with the oxygen to form oxides of carbon and 
water vapor, while the nitrogen is liberated in the free state. 
These explosives are all made by the action of a mixture of 



SOAP; GLYCERIN; EXPLOSIVES 



323 



nitric and sulfuric acid upon certain compounds, the name of 
the explosive ordinarily indicating the compound used. The 
most important of the nitro-explosives are given below : 




Fig. 189. Powder grains for large guns (natural size) 

(a) Nitrocellulose. Approximate composition, C 6 H 7 o (N0 ) 2 . 
This substance (p. 246) is a far more powerful explosive than 
gunpowder. If ignited, it will, under ordinary conditions, burn 
quietly. If subjected to a sudden shock 
(such as may be produced by the explo- 
sion of a small percussion primer), the 
nitrocellulose decomposes with enor- 
mous violence. The products of the 
decomposition are all colorless gases ; 
hence the use of this explosive in 
making smokeless gunpowder. When 
used for this purpose it is necessary to 
modify the pure material somewhat, 
as otherwise the violence of the explo- 
sion would shatter an}^ firearms in 
which the powder was used. This is 
done by mixing nitrocellulose with 
sufficient alcohol and ether to form a 
jelly. This is then molded into the 
form of rods (grains), with a number 
of perforations through the rods. The size of the grains varies 
with the size of the guns in which the powder is used. Fig. 189 
shows the form of the grains used in the large guns of our navy. 




Fig. 190. A stick of dyna- 
mite -with fuse attached 



324 



CHEMISTRY AND ITS USES 



(b) Nitroglycerin (C 8 H 5 (N0 8 )g). This compound (p. 322) resem- 
bles nitrocellulose in the violence of its explosive effects. The 
changes taking place in its decomposition are represented in a 
general way by the following equation : 

4 C 3 H 6 (N0 8 ) 3 —J- 12 C0 2 + 6 N 2 + 10 H 2 + 2 

One volume of nitroglycerin yields on explosion about 1300 vol- 
umes of gas, which is expanded by the heat of the reaction to 
over 10,000 volumes. Pure nitroglycerin is very dangerous because 




Fig. 191. Rows of depth bombs (A, A, A, A) on the rear of a 
submarine destroyer 

It was these destroyers and bombs which overcame the submarine 
menace in the World War 

of the ease with which it is set off. Large quantities are used in 
making dynamite, in which form it is not exploded so readily by 
jarring and can be transported with less danger. Ordinary dyna- 
mite (Fig. 190) consists of a mixture of sodium nitrate, nitro- 
glycerin, and wood pulp, the latter acting like a sponge to absorb 
the nitroglycerin. Gelatin dynamite consists of nitrocellulose and 
nitroglycerin mixed together to form a jelly. It is a very powerful 
explosive, since it contains, no inert material, 



SOAP ; GLYCERIN ; EXPLOSIVES 325 

(c) Trinitrotoluene (C 7 H 5 (N0 2 ) 3 ). This explosive, commonly 
known as T.N.T., is a white solid, but colors on exposure to air. 
It is too powerful to be used in guns, but is especially adapted for 
use in bombs and torpedoes. It was largely used in the World 
War. Eig. 191 shows two rows of depth bombs (A, A) on the rear 
of a submarine destroyer. Each bomb contains about 300 pounds 
of T.N.T. If a submarine was known to be in a certain locality, 
these bombs were dropped into the water at that point, from the 
stern of the rapidly moving destroyer, and the pressure of the water 
caused them to explode when they reached a certain depth. 

(d) Picric acid or trinitrophenol (C 6 H 2 OH(N0 2 ) 3 ). This is a yel- 
low solid (p. 234) and was the favorite explosive of the French in 
the World War. Large quantities were made in the United States 
and shipped to France. 

EXERCISES 

1. Illustrate the meaning of the term by-product, using the soap 
industry as an example. 

2. Could calcium hydroxide be used as the alkali in soap manu- 
facture ? 

3. Write the equation for the reaction (a) between olein and sodium 
hydroxide ; (6) between palmitin and potassium hydroxide. 

4. (a) To what class of substances (acid, base, or salt) does soap 
belong? (b) Account for the fact that a solution of soap reacts basic 
(p. 308). 

5. Can you suggest any reason for the cleansing properties (a) of 
aqua ammonia? (b) of borax? (c) of sodium carbonate? 

6. Should you expect soap to lather in rain (cistern) water ? 

7. Do you see any reason for adding naphtha to soap as is done in 
making naphtha soaps? 

8. Families used to make their soap by heating the fat refuse of 
foods with a solution prepared by extracting wood ashes with water, 
(a) What was the alkali used (p. 311) ? (b) What kind of soap resulted ? 

9. What effect does the softening of a city water supply have upon 
soap consumption ? 

10. Give the steps in making candles from garbage. 

11. Sodium nitrate is much less expensive than potassium nitrate. 
Why not use it in making gunpowder? 



326 CHEMISTRY AND ITS USES 

12. Why are the nitro-explosives more powerful than gunpowder? 

13. Why are some powders smokeless and others not? 

14. (a) Has T.N.T. sufficient oxygen to burn completely the carbon 
and hydrogen present? (b) Account for the fact that when T.N.T. 
explodes, great volumes of black smoke are formed. 

15. Should you expect the following explosives to be smokeless: 

(a) nitroglycerin? (&) picric acid? 

16. (a) Why did not the Germans use nitroglycerin explosives dur- 
ing the World War ? (&) Where did they obtain the nitric acid necessary 
for making explosives ? 

17. Calculate the percentage of nitrogen (a) in nitroglycerin ; 

(b) in T.N.T. 



CHAPTER XXXVI 
THE CALCIUM FAMILY 



Name 


Symbol, 


Atomic 
Weight 


Density 


Formula of 
Chloride 


Calcium 

Strontium 

Barium 


Ca 
Sr 
Ba 


40.07 

87.63 

137.37 


1.55 
2.54 
3.75 


CaCl 9 
SrCl 2 
BaCl 2 



The family. The calcium family consists of the very 
abundant metal calcium and the rarer metals, strontium 
and barium. Like the alkali metals, they are acted upon by 
both water and air, and on this account do not occur in a 
free state in nature. They are bivalent, so that the formulas 
of their salts differ from the formulas of the corresponding 
salts of the alkali metals. 



Calcium 

Introduction. The metal calcium, like sodium and potassium, 
is little used, but its compounds are numerous and abundant. 
Its most familiar compounds are calcium carbonate in the form 
of limestone and marble, and calcium phosphate, which occurs 
abundantly in nature and constitutes the mineral constituents 
of our bones. The metal itself is silver-white and light in 
weight but quite hard. It is prepared by the electrolysis of 
its melted chloride. It burns in oxygen with dazzling bril- 
liancy and, like sodium, decomposes water, forming calcium 
hydroxide and hydrogen. 

327 



328 



CHEMISTRY AND ITS USES 



Calcium carbonate (CaC0 3 ). Enormous quantities of calcium 
carbonate occur in nature. Limestone is the most abundant 
form and sometimes constitutes whole mountain ranges. 

Limestone is never pure calcium 
carbonate, but contains vari- 
able percentages of magnesium 
carbonate, clay, silica, and com- 
pounds of iron. Pearls, coral, 
and various shells are largely 
calcium carbonate. Calcite is a 
very pure, crystalline form and 
often is found in large transpar- 
ent crystals (Fig. 192) called 
Fig. 192. A crystal of Iceland spar Iceland spar. Marble is composed 

of very small calcite crystals. 
In the laboratory pure calcium carbonate can be prepared 
by mixing solutions of a calcium salt and some carbonate : 




Na CO„ + CaCl, 



CaC0 o + 2 NaCl 



When prepared in this way it is a soft white powder often 
called precipitated chalk, and is much used as a polishing 
powder (tooth powder). 

The natural varieties of calcium carbonate find many uses, 
as in the preparation of lime and of carbon dioxide ; in metal- 
lurgical operations, especially in making iron and steel; in 
the manufacture of soda and glass ; for building-stone and 
as ballast for roads. 

Calcium acid carbonate (Ca(HC0 3 ) 2 ). Calcium carbonate is 

almost insoluble in pure water. It readily dissolves, however, 

in water which holds carbon dioxide in solution (p. 113). This 

is due to the fact that the carbonate combines with the carbonic 

acid present in the water to form calcium acid carbonate, which 

is soluble : 

CaC0 3 + H 2 C0 3 ^=tCa(HC0 3 ) 2 



THE CALCIUM FAMILY 



329 



The resulting acid carbonate exists only in solution, since it 
is unstable and decomposes into the normal carbonate on 
heating or on evaporation of its solution. 

Formation of caves. Natural waters always contain more or less 
carbon dioxide in solution. In the case of certain underground 
waters the amount of carbon diox- 
ide is comparatively large, being 
held in solution by pressure. Such 
waters have a marked solvent 
action upon limestone, dissolving 
both the calcium carbonate and 
the magnesium carbonate. In 
certain localities this solvent ac- 
tion, continued through geologi- 
cal ages, has resulted in the 
formation in limestone rock of 
large caves, such as the Mam- 
moth Cave in Kentucky. Water 
trickling through the roofs of 
these caves evaporates, leaving 
a deposit of calcium carbonate, 
which, as the process continues, 
often forms icicle-shaped masses 

known as stalactites ; or the water may drip upon the floor of the 
cave, forming similar masses known as stalagmites (Fig. 193). 

Commercial methods for softening water. Ordinary hard 
waters, such as well and river waters, contain not only the 
acid carbonates of magnesium and calcium but also the 
chlorides and sulfates of these metals, together with small 
quantities of other salts. Experiments have shown that such 
waters may be softened by the addition of calcium hydroxide 
and sodium carbonate (Fig. 194). The calcium hydroxide 
converts the acid carbonates of calcium and magnesium into 
the normal carbonates, which, being insoluble, settle out: 

Ca(HC0 3 ) 2 + Ca(OH) 2 — y 2 CaC0 3 + 2 H 2 




Fig. 193. Stalactites and 
stalagmites 



330 



CHEMISTRY AND ITS USES 



The sodium carbonate used converts the chlorides and sulfates 
of calcium and magnesium into the insoluble carbonates : 

CaS0 4 + Na„C0 3 — >■ CaC0 3 + Na 2 S0 4 

Water softened in this way contains sodium sulfate and 
chloride, but these salts are not objectionable. 




Fig. 194. Softening the water supplied to Columbus, Ohio 

In softening the water supply of a city the calcium hydroxide and sodium 

carbonate are added to the water and mixed thoroughly, as shown in the figure. 

The mixture is then run into large settling basins where the solids settle. The 

clear, soft water is then filtered and run into the city mains 

Calcium oxide (lime) (CaO). Pure calcium oxide can be 
prepared by burning calcium or by heating its carbonate. 
The more or less impure oxide (commercially known as lime 
or quicklime^) is always prepared by strongly heating limestone 
in large furnaces called limekilns: 

CaCQ, ^CaO + C0 2 

Lime, when pure, is a white amorphous substance. It melts 
only in the heat of the electric furnace. Water acts upon 
lime with the evolution of a great deal of heat, — whence 



THE CALCIUM FAMILY 



331 



the name quicklime, or live lime, — the process being called 
slaking. The equation is 

CaO + H 2 >■ Ca(OH) 2 + 15,540 cal. 

Because of its affinity for water it is used to dry gases. It 
also absorbs carbon dioxide, forming the carbonate. Lime 
exposed to air is therefore gradually con- 
verted into the hydroxide and then into 
the carbonate, and will no longer slake 
with water. It is then said to be air-slaked. 
Lime is produced in enormous quantities 
for use in making calcium hydroxide. 

Commercial production of lime. A vertical 
section of the newer form of limekiln is 
shown in Fig. 195. The kiln is about 50 ft. 
in height. A number of fire boxes, or fur- 
naces A, A, are built around the lower part, 
all leading into the central stack. The kiln 
is filled with limestone through a swinging 
door B. The hot products of combustion are 
drawn up through the kiln, and the limestone 
is gradually decomposed by the heat. The 
bottom of the furnace is so constructed that 
a current of air is drawn in at C, and this 

serves the double purpose of cooling the hot lime at the base of 
the furnace and of furnishing heated oxygen for the combustion. 
The lime is dropped into cars run under the furnace. Ordinarily a 
number of these kilns are operated together, as shown in Fig. 196. 

Calcium hydroxide (slaked lime, hydrated lime) (Ca(OH) 2 ). 

This compound is prepared by adding water to lime, as ex- 
plained above. When pure it is a light white powder. It 
is sparingly soluble in water, forming a solution called lime- 
water. Owing to its cheapness it is used in the industries 
whenever an alkali is desired. It is used in the prepara- 
tion of ammonia, bleaching powder, and the hydroxides of 




Fig. 195. A vertical 

section of a modern 

limekiln 



332 



CHEMISTRY AND ITS USES 




sodium and potassium. It is also used to remove sulfur com- 
pounds and carbon dioxide from coal gas, to remove the hair 
from hides in making leather, for making mortar and plaster, 

and for neutralizing 
the acids in soils 
(p. 335). 

Mortar and plaster. 

Mortar is a mixture 
of calcium hydroxide 
and sand. When it is 
exposed to the air or 
spread upon porous 
materials, moisture 
is removed from it 
(partly by absorption 
into the porous ma- 
terials and partly by 
evaporation) and the 
mortar becomes firm, 

or sets. At the same time carbon dioxide is slowly absorbed from 

the air and hard calcium carbonate is formed : 

Ca(OH) 2 + C0 2 ^CaC0 3 + H 2 

By this combined action the mortar becomes very hard and adheres 
firmly to the surface upon which it is spread. The sand serves 
to give body to the mortar and makes it porous. It also prevents 
too much shrinkage. Plaster is a mixture of calcium hydroxide 
and hair or wood fiber, the latter serving to hold the mass together. 

Bleaching powder. When chlorine is passed over calcium 
hydroxide there is formed a white solid compound having 
the formula CaOCl 2 and known as bleaching powder, chloride 
of lime, or simply bleach : 

Ca(OH) 2 + Cl 2 — >■ CaOCl 2 + H 2 

When this compound is treated with an acid, chlorine is 
evolved : 



Fig. 196. A group of limekilns in a modern plant 






CaOCl +HSO, 



CaSO ,+H.O+Cl. 



THE CALCIUM FAMILY 333 

When exposed to the air bleaching powder is slowly acted upon 
by moisture and carbon dioxide, evolving hypochlorous acid, 
which is a good disinfectant. Bleaching powder is prepared 
in large quantities for use as a bleaching agent, as a disin- 
fectant, and as an agent for purifying city water-supplies. 

In the commercial preparation of bleaching powder the calcium 
hydroxide is spread to a depth of 2 or 3 inches upon the floor of 
a room, usually made of lead or concrete (Fig. 197). The chlorine 
enters near the top at A. Any unabsorbed chlorine passes out at 



WMMMMMMM#MMMM/MM/MM/mm t 




B 




Fig. 197. Chambers for the manufacture of bleaching powder 

B and into the adjoining chamber at C. The commercial product 
ordinarily contains about 35 per cent of chlorine. It is shipped 
in sealed packages or iron drums. 

Calcium sulfate (CaSOJ. Calcium sulfate occurs in nature 
in several different forms, the most common of which is 
gypsum (CaS0 4 • 2 H 2 0). This is quarried in large amounts 
in New York, Iowa, and Michigan. It is used as a filler in 
making paper (p. 250), as a constituent of fertilizers, and 
especially in making plaster of Paris. It is but slightly soluble 
in water. 

Plaster of Paris ((CaS0 4 ) 2 . H 2 0). This is a fine white 
powder obtained by carefully heating gypsum. When water 
is added to the powder a plastic mass is formed which quickly 
hardens, or sets. This property makes it valuable for molding 
casts, for stucco work, and for a finishing coat on plastered 
walls. Broken bones are often held in place by casts of 
plaster of Paris until they grow together. 



334 



CHEMISTRY AND ITS USES 



Calcium carbide (CaC 2 ). This compound is prepared on a 
large scale for use in the manufacture of acetylene (p. 215) 
and in making fertilizers. It is made by heating a mixture of 
lime and coke in an electric furnace : 



CaO + 3 C 



CaO + CO 



- 




Fig. 198. A furnace for making calcium carbide 



The pure carbide is a colorless, transparent solid. The com 
mercial article is a dull-gray porous substance which contains 

many impurities. It is 
placed on the market 
in air-tight containers. 



Commercial prepara- 
tion. While calcium car- 
bide was first prepared 
in 1836, it was not un- 
til 1893 that it became 
a commercial product. 
The general principles 
involved in its prepara- 
tion are illustrated in Fig. 198. A large brick room open at the 
top is fitted with carbon electrodes A, A, A and filled with a mix- 
ture, B, of lime and coke. Sufficient current is used to secure a 
temperature of about 2000°. The carbide melts as fast as it forms, 
C, C, and is drawn off from time to time at D. 

Calcium cyanamide (CaCN 2 ). When nitrogen is passed over 
hot calcium carbide the two react, forming a compound known 
as calcium cyanamide : 

CaC 2 + N 2 t-> CaCN 2 + C 

The commercial product is impure, containing about 60 per 
cent of the calcium cyanamide; the impurities are lime and 
carbon. This is ground, mixed with water (which slakes 
the lime present), and in this form sold as a fertilizer under 
the name of cyanamide. Its value as a fertilizer lies in the 
fact that all its nitrogen is available as a plant food. 




Fig. 199. View in a plant for making calcium carbide, showing the 
large carbon electrodes which dip into the furnace below 




Fig. 200. View in a plant for making calcium cyanamide 

Calcium carbide is placed in the perforated steel cans shown in the back- 
ground. These are then lowered into the furnaces shown in the foreground. 
The carbide is heated to about 1000°, and then pure nitrogen obtained from 
liquid air is passed over it, the two combining to form the cyanamide 



AIR PRODUCTS 






s 



AIR 



NITROGEN 



1 



OXYGEN 

(OXY-AGETYLENE WELPli 
(MEDICINAL) , 



CALCIUM CYANAMIDE 
.; fertilizer'! 



I 



PHOSGENE 

/ WAR GAS AND \ 
\DYE INTERMEDIATE^ 



I 



AMMONIA 



1 



AMMONIUM NITRATE 

HIGH EXPLOSIVE) 
(FERTILIZER) 



1SF 




| 
NlTROtfS- OXIDE NITROeHCERIN CHLORPtCRIN CELLULOSE NITRATE 

(ANAESTHETiC) (EXPLOSlVF' (WAR GAS) (SMOKELESS POWDER 

(MEDICINAL) 



Courtesy of Science Service 



Fig. 201. Some products obtained from the air during the World War 



THE CALCIUM FAMILY 335 

Calcium cyanamide is also used as a source of ammonia as 
well as sodium cyanide, thus : 

CaCN 2 + 3 H 2 — y CaC0 3 + 2 NH 3 
CaCN 2 + C + 2 NaCl h CaCl 2 + 2 NaCN 

During the World War the government constructed a large 
plant (costing nearly $100,000,000) at Muscle Shoals, -Alabama, 
for preparing ammonia by this process. A part of the ammonia 
was oxidized to nitric acid (p. 169), and this was then combined 
with the unchanged ammonia to form ammonium nitrate, which 
was valuable as an explosive. The plant had a capacity of 300 tons 
of the nitrate daily. The chart on the opposite page shows some of 
the products formed from the atmosphere during the war. 

Calcium phosphate (Ca 3 (POJ 2 ). This important substance oc- 
curs in nature as the mineral phosphorite and as a constituent of 
apatite. Large amounts of it occur in the form of rock phosphate, 
which is found especially in Florida, Tennessee, and some of 
the Western states. It is also the chief mineral constituent of 
bones. Bone ash is, therefore, nearly pure calcium phosphate. 

Other compounds of calcium. Calcium chloride (CaCl 2 ) occurs in 
sea water and is formed in large quantities as a by-product in the 
Solvay process for making sodium carbonate. The anhydrous salt 
readily absorbs moisture and is used as an agent for drying gases. 
A solution of the salt is used as a brine in the manufacture of 
ice (p. 166). It has also been used to lay the dust on roads, and 
mines have been sprinkled with it in the hope of preventing dust 
explosions. Calcium acid sulfite (Ca(HS0 3 ) 2 ) is used in large quan- 
tities in the manufacture of paper (p. 250). A number of calcium 
silicates are known, and derive their chief interest from the fact 
that they are important constituents of cement and glass. 

Strontium and Barium 

General. These elements are much rarer than calcium, are diffi- 
cult to prepare, and have no commercial uses. Their compounds 
resemble those of calcium in composition and properties. Stron- 
tium compounds, especially the nitrate, when ignited with oxidiz- 
able substances, give a brilliant crimson color and on this account 






The hides of animals 
from which leather is 
made are treated with 
salt and stored away in 
large cellars until they 
are needed for tanning, 
as shown in A 



The hides are converted 
into leather hy the action 
of tannin (a complex 
substance present in the 
bark of certain trees, es- 
pecially the oak) or some 
substance which acts in 
a similar way. In B are 
shown the supplies of oak 
bark stored in the yard 
of a tannery 



Before treatment with 
tannin, the hides are first 
soaked in a strong solu- 
tion of calcium hydrox- 
ide (C) to loosen the 
hair, which is then easily 
removed 



Fig. 202, A, B, and C. The tanning of leather 



After the hair is removed 

the hides are scraped 

to remove all adhering 

flesh, as shown in D 



As shown in E, the hides 
are then placed in tan- 
ning vats for a number 
of weeks. These contain 
water and ground tan- 
bark, the tannin of which 
slowly changes the hides 
into the tough, raw or 
unfinished leather 



In F the raw leather is 
being washed, scoured, 
ironed, greased, and 
dyed (if a colored leather 
is desired). The particu- 
lar treatment of the raw 
leather depends upon the 
kind of leather desired 




Fig. 202, D, E, and F. The tanning of leather 






336 CHEMISTRY AND ITS USES 

are used in the manufacture of red lights. Under similar condi- 
tions barium nitrate gives a green light. The following compounds 
of barium are of importance : 

Barium chloride (BaCl 2 '• 2 H 2 0). Barium chloride is a white solid. 
It is used in the laboratory as a reagent to detect the presence of 
sulfuric acid or soluble sulfates, reacting with these to form the 
insoluble barium sulfate. 

Barium sulfate (barite) (BaS0 4 ). Barium sulfate occurs in nature 
as a heavy white mineral known as barite or heavy spar. It is 
precipitated as a crystalline powder when a barium salt is added 
to a solution of a sulfate or to sulfuric acid : 

BaCl 2 + H 2 S0 4 >■ BaS0 4 + 2 HC1 

It is used in large quantities in the manufacture of paints. 

The tanning of leather. The use of calcium hydroxide in the 
tanning of leather suggests the brief description of the process 
given in connection with Fig. 202, A, B, C, D, E, and F. 

EXERCISES 

1. (a) What properties do the members of the calcium family have 
in common with the alkali metals ? (b) In what respects do they differ ? 

2. Do the members of the calcium family show any gradation of prop- 
erties corresponding to their atomic weights ? (See table at head of chapter.) 

3. What simple test will enable you to tell whether a rock contains 
a carbonate ? 

4. Some samples of Iceland spar and gypsum look very much alike. 
How could you distinguish between them ? 

5. Beads made of celluloid (p. 246) are often dyed and sold as coral 
beads. Can you suggest any simple way of detecting the fraud? 

6. If you blow exhaled air through limewater, the clear solution 
becomes cloudy, but as you continue, it clears again. If now you boil 
the clear solution it again becomes cloudy. Explain. 

7.- Account for the deposits in teakettles in which hard waters are 
boiled. 

8. (a) Is the reaction represented by the equation 

CaC0 3 >-CaO + C0 2 

reversible ? (b) If so, how can you make it go in either direction ? 






THE CALCIUM FAMILY 337 

9. (a) Would lime after long exposure to the air serve for making 
mortar? (b) Explain. 

10. Write the equation for the reaction that takes place when calcium 
hydroxide is used for liberating the ammonia from ammonium chloride. 

11. Enormous quantities of calcium chloride are formed as a by- 
product in the Solvay process for the manufacture of sodium carbonate 
(p. 308). Account for its formation. 

12. (a) Show how the nitrogen of the atmosphere can be utilized in 
preparing ammonia through the agency of calcium carbide, (b) In 
what other way can the same result be obtained (p. 165) ? 

13. Mention the advantages gained by a city from softening its 
water supply. 

14. Why do calcium and magnesium salts make a water hard, while 
sodium salts do not? 

15. What term do we apply to compounds like calcium chloride that 
become moist on exposure to air (p. 314) ? 

16. If you used large quantities of bleaching powder, would you be 
willing to base the price solely on the weight of the product furnished ? 

17. In making calcium cyanamide very high temperatures and very 
low temperatures are utilized. In what connection is each of these 
temperatures employed ? 

18. What weight of limestone is burned in a plant which has an 
output of 25 tons of lime daily ? 

19. (a) What weight of water is required for slaking 1 kg. of lime? 
(&) How many calories of heat will be evolved in the slaking ? 

20. The water used by a certain city contains 120 g. of calcium acid 
carbonate in 100 gal. of water, (a) What weight of calcium hydroxide 
is necessary to soften each 100 gal. of the water (see equation, p. 329) ? 
(b) Would it make any difference if you added more than this weight 
of the calcium hydroxide? 

21. A certain factory produces 50 tons of bleaching powder daily. 
How many pounds of calcium hydroxide are required daily ? 

22. (a) If bleaching powder is pure CaOCl 2 , what per cent of chlo- 
rine would it contain ? (b) Compare your result with the amount in the 
commercial product (p. 333). 

23. What weight of gypsum is required in making 1000 kg. of plaster 
of Paris? 



CHAPTER XXXVII 



SOILS AND FERTILIZERS 



Formation of soils. Soil results from disintegration of 
rock. Many agencies contribute to this change, prominent 
among them being the action of water and air. The glaciers 

which once covered a 
large portion of the 
globe exerted a very 
powerful grinding ef- 
fect, breaking off large 
bowlders and reducing 
them to a powder. 
Water acts in a num- 
ber of ways. It finds 
its way into the cracks 
and crevices, and on 
freezing expands and 
so tends to break the 
rocks ; it dissolves out 
certain rock constitu- 
ents so that the re- 
mainder of the rock 
tends to crumble ; and 
running water carry- 
ing particles in suspen- 
sion grinds the rocks 
with which it comes in contact. Air also acts upon some 
rocks chemically, while strong winds carrying sand dust exert 
a powerful sand-blast effect. When plants get a foothold 

338 




Fig. 203. The formation of soil through the 
disintegration of rocks 



SOILS AXD FEBTILIZEKS 



339 



they assist greatly, not only by the mechanical action of their 
roots but also through the chemical action of the various 
products formed in their decay. The soil is the result of all 
these agencies acting through long periods of time (Fig. 203). 
Different kinds of soil. Since the rocks vary widely in com- 
position, it follows that the resulting soils will vary in a cor- 
responding way. Thus we have limestone soil, formed chiefly 




Fig. 204. Neutralizing acidity in soil by an application of calcium 
hydroxide or ground limestone 

by the weathering of limestone ; and sandy and clay soils, 
which are richer in silicates. When plants decay certain acid 
products are formed, so that soils tend to become acid. Most 
crops do not thrive on such soil. If the soil contains lime- 
stone, the acids will be neutralized as fast as formed and the 
soil thus kept sweet ; if the soil does not contain limestone, 
then ground limestone or calcium hydroxide is added to the 
soil (Fig. 204). 

Soils differ not only in composition but also in their physi- 
cal properties. Thus, some soils are light and easily allow 
water and air to circulate through them (these conditions are 
essential to a fertile soil), while others are compact and 



340 



CHEMISTRY AND ITS USES 



impervious to water. Crops do not thrive on wet soils ; hence 
the necessity of drainage (Fig. 205). A good soil must have 
many properties of a colloid jelly, in which adsorption of salts 
and gases plays an important part in holding plant food. 

Plant food and fertilizers. 
With the exception of carbon 
dioxide (and possibly a little 
oxygen) absorbed from the air, 
the growing plant derives its 
nourishment from the soil. In 
order that vegetation may 
thrive it is essential, therefore, 
that the soil should contain 
an adequate supply of appro- 
priate plant food, and this in- 
cludes both mineral matter 
and organic matter (humus'). 
Moreover, since this supply is 
continually being drawn upon 
by the growing plant, it is 
necessary, in order that the 
soil may retain its fertility, 
that the ingredients so with- 
drawn shall be returned to it. 
It is for this purpose that 
fertilizers are used. 
Constituents of fertilizers. While a number of elements are 
essential to the growth of the plant, experience has shown 
that in general the fertility of a soil that has the necessary 
physical properties may be maintained by adding three sub- 
stances: (1) nitrogenous matter, (2) phosphates of calcium, 
and (3) compounds of potassium. It seems probable that the 
element sulfur also plays an important part in the growth of 
plants, although its action is not thoroughly understood. 



,. t : mgj 


-. * ; ' r* /^>v^i 


I^^|^^'.:. 








::■"•■.'' 


^ 


.. 




'■ Mm M I W I 
If ' " W -* 1 

mm 
i \ 


'mm' ' 


£ ' 


. * || 


w H ■ 



Fig. 205. Draining wet soils so as 
to make them fertile 



SOILS AND FEKTILIZERS 341 

Sources of fertilizers. The commercial sources of each of 
the constituents of fertilizers are as follows : 

1. Nitrogenous matter. This is obtained from a number of 
sources : sodium nitrate, ammonium sulfate, and cyanamide ; 
also nitrogenous organic matter, such as dried blood, the waste 
from slaughterhouses, and, especially, animal excrements. 

2. Phosphates. Ground bones are especially valuable, since 
they contain some nitrogen in addition to calcium phosphate. 
This source, however, is entirely inadequate, and the great 




Fig. 206. Mining phosphate rock in Florida 

supply comes from the rock phosphates, which contain about 
70 per cent calcium phosphate. These rock phosphates are 
quarried in large quantities, especially in Florida (Fig. 206) 
and Tennessee. Since calcium phosphate is nearly insoluble, 
the rock is ground and then treated with sulfuric acid. This 
converts the insoluble calcium phosphate into the soluble 
calcium acid phosphate (CaH 4 (P0 4 ) 2 ) : 

Ca 8 (P0 t ) 2 + 2 H 2 S0 4 —J- 2 CaS0 4 + CaH 4 (P0 4 ) 2 

The calcium sulfate also adds to the value of the fertilizer, 
furnishing sulfur and improving the physical qualities of the 
soil. Certain products (slags) formed in the manufacture of 
steel contain phosphorus and are also used in fertilizers. 



342 



CHEMISTRY AND ITS USES 



3. Potassium compounds. These are obtained principally from 
the natural deposits in Germany and France (Fig. 184). Wood 
ashes are excellent, but the supply is limited. 

Commercial fertilizers. As a rule the fertilizers on the 
market are mixtures of the three fundamental materials re- 
ferred to above. The composition is varied according to the 
crop to be grown as well as to the nature of the soil; for 
example, potatoes demand a fertilizer rich in potassium, while 

a cereal, such as wheat, is 
benefited more by one rich in 
phosphates. Instead of using 
a fertilizer containing all three 
constituents, it is preferable to 
find out by experiment just 
what plant food is lacking in 
individual soils and then to 
make a proper mixture of such 
fertilizing materials as will fur- 
nish the desired food. Fig. 207 
shows the result of such an 
experiment. Pot 1 shows the 
result obtained with no fertilizer, pot 2 the effect of one 
combination, and pot 3 that of another. 

Utilization of atmospheric nitrogen. It has been pointed 
out that with few exceptions plants have not the power of 
assimilating free nitrogen (p. 79). Moreover, it is inevitable 
that the supplies of sodium nitrate and ammonium sulfate, 
which are now the chief nitrogenous products used in the 
manufacture of commercial fertilizers, will sooner or later 
become exhausted. It has been found possible to utilize the 
inexhaustible supply of nitrogen in the atmosphere by con- 
verting the nitrogen into compounds which contain the ele- 
ment in a form available for plant food. We have already 
discussed these methods (pp. 169, 334). 









1 




Br/ / 




.-'""H 








wx 


1 f i ' 






lii 




% . 


• 


Jmi 














Mj 




^ 





Fig. 207. The effect of fertilizers 
upon the growth of plants 



SOILS AND FERTILIZERS 343 

EXERCISES 

1. Give examples of soil formation that have come under your 
observation. 

2. Have you ever seen any evidence that the roots of plants and 
trees exert a strong mechanical pressure ? 

3. (a) Suppose the crops grown on a soil were not removed. What 
result would this have on the fertility of the soil ? (b) Would the results 
be influenced by the nature of the crop ? 

4. Give two different methods for changing the nitrogen of the air 
into fertilizers. 

5. Suggest a possible way of utilizing the large plant built at Muscle 
Shoals (p. 335) during the war for the manufacture of explosives. 

6. AVhat kind of crops enrich the soil (p. 79) ? 

7. What compound is present. in wood ashes that makes the ashes 
a valuable fertilizer? 

8. Would wood ashes, if available, be effective to "sweeten" soils? 

9. Would air-slaked lime (p. 331) serve to sweeten soils? 

10. Ground bones are sometimes used as fertilizers, but their action 
is slow. Explain. 

11. How could you find out what mineral products are withdrawn 
from the soil by different kinds of crops? 

12. In making fertilizers the sulfuric acid used contains about 50 per 
cent of hydrogen sulfate. W^hat weight of such an acid is necessary for 
the treatment of 1 ton of rock phosphate containing 70 per cent of 
calcium phosphate ? 



CHAPTER XXXVIII 
THE MAGNESIUM FAMILY 



Name 


Symbol 


Atomic 
Weight 


Density 


Melting 
Point 


Boiling 
Point 


Oxide 


Magnesium . . 

Zinc 

Cadmium . . . 


Mg 
Zn 
Cd 


24.32 
65.37 
112.4 


1.74 
7.10 

8.64 


651° 

419.4° 

320.9° 


920° 
950° 

778° 


MgO 
ZnO 
CdO 



The family. In the magnesium family are included the 
four elements magnesium, zinc, cadmium, and mercury. 
Between the first three of these metals there is a close family 
resemblance. In some respects mercury resembles copper more 
closely and will be studied in connection with that metal. 

Magnesium 

General. Magnesium is one of the abundant elements, 
although it never occurs native. Like sodium and calcium 
it is obtained pure by the electrolysis of its compounds, the 
mineral carnallite (p. 311) being used for this purpose. The 
pure metal is silver white and so light that it will just sink 
in water. Air tarnishes it somewhat. The common acids 
readily dissolve it. When it is ignited it burns with great 
brilliancy, even in air, and gives a light rich in the rays 
which act upon photographic plates; hence its use for making 
flash lights and fireworks and for the flares used in lighting 
battlefields at night. Ordinary flash-light powder is a mixture 
of powdered magnesium and potassium chlorate or some similar 
oxidizing agent. Magnesium is also used in making alloys. 

344 



THE MAGNESIUM FAMILY 345 

Occurrence. The most common forms of rocks containing 
magnesium are dolomite (CaC0 3 • MgC0 3 ) and magnetite 
(MgC0 3 ). Asbestos, talc, and serpentine are silicates of 
magnesium. 

Magnesium oxide (MgO). This is the soft white powder 
often known as magnesia. It resembles lime in many respects, 
but is even more infusible ; hence its use for making fire 
brick for lining furnaces. It combines with water to form 
the hydroxide Mg(OH) 2 , which resembles calcium hydroxide. 

Magnesium chloride. Magnesium chloride occurs in many 
natural waters along with the chlorides of sodium and calcium. 
It is coming into use for making floors. A common floor cover- 
ing is made by adding magnesium chloride to a mixture of 
wood fiber or asbestos and magnesium oxide. With the proper 
amount of water this mixture sets, forming a tough, some- 
what elastic mass. 

Magnesium carbonate (MgC0 3 ). The pure carbonate occurs 
as magnesite. More often it is associated with calcium car- 
bonate to form dolomite or dolomitic limestone, which forms 
whole mountain ranges and strata in the earth's crust. Water 
containing carbon dioxide acts upon magnesium carbonate, 
just as it acts upon calcium carbonate (p. 328). 

Magnesium sulfate (MgS0 4 ). Magnesium sulfate differs from 
calcium sulfate in that it is very soluble in water. It occurs 
in some springs and crystallizes in the form MgS0 4 • 7 H 2 0. 
This hydrate is used in medicine under the name of Epsom 
salt Large beds of the hydrate occur in Wyoming and 
Washington. It has a number of commercial uses, especially 
in the dye industry. 

Magnesium silicates. Asbestos (CaMg 3 (Si0 4 ) 2 ) occurs in 
quantities in the province of Quebec, Canada, and is used 
for making fireproof shingles, board, and for covering pipes 
to diminish heat radiation. Talc (Mg 3 H 2 (Si0 3 ) 4 ), also known 
as soapstone, is valuable for making sinks and table tops and, 



346 



CHEMISTRY AND ITS USES 



in finely ground form (French chalk), for toilet powders. 
Serpentine is a greenish silicate of some use as a building 
stone, while meerschaum, also a silicate, is used for pipe 
bowls and similar articles. 

Boiler scale. When water which contains certain salts in solu- 
tion is evaporated in steam boilers, a hard, insoluble material 

called scale deposits in the 
boiler. The formation of this 
scale may be due to several 
distinct causes : 

1. To the deposit of calcium 
sulfate. This salt, although 
sparingly soluble in cold 
water, is almost completely in- 
soluble in superheated water. 
Consequently, when water 
containing it is heated in a 
boiler the calcium sulfate pre- 
cipitates and forms a hard 
scale. 

2. To decomposition of acid 
carbonates. As we have seen, 
calcium and magnesium acid 

carbonates are decomposed on heating, forming insoluble normal 
carbonates : Ca(H C0 3 ) 2 ^ CaC0 3 + H 2 + CO a 

3. To hydrolysis of magnesium salts. Magnesium chloride and, to 
some extent, magnesium sulfate undergo hydrolysis when super- 
heated in solution, and the magnesium hydroxide, being sparingly 
soluble, precipitates : 

MgCl 2 -f 2 H 2 y Mg(OH) 2 + 2 HC1 

This scale adheres tightly to the boiler tubes in compact layers 
(Fig. 208) and, being a nonconductor of heat, causes much waste 
of fuel. It is very difficult to remove, owing to its hardness and 
its resistance to reagents. Thick scale sometimes cracks, and the 
water coming in contact with the overheated iron occasions an 
explosion. 




Fig. 208. Cross section of a boiler tube 
showing the deposit- of boiler scale 



THE MAGNESIUM FAMILY 347 

Zjnc 

Properties. Zinc is the first of the metals so far studied 
that is familiar to all of us. It is a bluish-white metal, about 
as heavy as iron. If heated strongly in the air it takes fire 
and burns to the oxide ZnO. If the metal is melted and 
allowed to cool, it crystallizes and at ordinary temperature is 
quite hard and brittle ; but if heated to between 100° and 150° 
it becomes malleable and can be rolled into thin sheets, and in 
this form the metal retains its softness and malleability at 
ordinary temperatures. When melted and poured into water 
it forms thin, brittle flakes known as granulated or mossy zinc ; 
it is this form which is used in preparing hydrogen. Air 
tarnishes it but exerts no further action. It dissolves in all 
the common acids. 

When the metal is quite pure, sulfuric and hydrochloric acids 
act upon it very slowly ; when, however, it contains small amounts 
of other metals, such as magnesium or copper, or when it is merely 
in contact with another metal, brisk action takes place and hydro- 
gen is evolved. For this reason, when pure zinc is used in the 
preparation of hydrogen a few drops of copper sulfate are often 
added to the solution to assist the chemical action. 

Occurrence an'd metallurgy. Zinc does not occur free in 
nature, neither is it a constituent of common rocks. Its 
chief ores are the following: sphalerite (ZnS), zincite (ZnO), 
smithsonite (ZnC0 3 ), and franldinite (ZnO • Fe 2 8 ). The 
chief zinc-producing states are Missouri, Montana, and 
New Jersey. 

To obtain the metal the ore is first roasted in the air, whereby 
the zinc in the ore is changed into the oxide. The oxide is then 
heated with carbon, which unites with the oxygen, liberating the 
zinc. The details of the process vary with the composition of the 
ore. Commercial zinc is often impure and contains small percent- 
ages of carbon, arsenic, and iron. 



348 CHEMISTRY AND ITS USES 

Uses of zinc. The chief use of zinc is in the manufacture 
of galvanized iron. This consists of sheets of iron covered with a 
thin layer of zinc, which protects the iron from rusting. About 
two thirds of all the zinc produced is used in this way. Sheet 
zinc is. used as a lining for sinks and water-containers. Large 
quantities of the metal are used in making brass and other 




Fig. 209. The manufacture of galvanized sheet iron 

alloys (p. 282), in the construction of electrical batteries, and 
in separating silver from lead. In the laboratory it is used in 
the preparation of hydrogen and, in the form of zinc dust, as 
a reducing agent. 

Manufacture of galvanized iron. Fig. 209 shows the method 
used in making galvanized iron. The plates of iron pass under 
the rollers at A and on into the pot of melted zinc B. The zinc 
adheres to the iron, and the resulting plate is passed under 
the roller C to remove the excess of zinc and to render the sur- 
face smooth. Sometimes the zinc is deposited on the iron by 
electrolytic methods. 



THE MAGNESIUM FAMILY 349 

Compounds of zinc. In general the compounds of zinc 
resemble those of magnesium in formula and appearance, 
although they differ markedly in chemical conduct. The most 
important of them are listed below: 

Zinc oxide (ZnO). This oxide occurs in impure state in 
nature. Commercially it is generally made by burning the 
metal in air. It is a pure-white powder which, under the 
name of zinc white, is much used as a white pigment in paints. 
It has an advantage over white lead in that it is not changed 
in color by sulfur compounds (which are likely to be present 
in the air of manufacturing districts), while lead compounds 
turn black. Many thousand tons of zinc oxide are used in 
paints each year. It is also used as a filler in the manufac- 
ture of rubber goods, especially automobile tires. 

Zinc sulfate (ZnSOJ. This salt is readily crystallized from 
concentrated solutions in transparent colorless crystals which 
have the formula ZnSO„ • 7 H„0 and are called white vitriol. 
It is used in dyeing and in medicine. 

Zinc chloride (ZnCl 2 ). This salt is very soluble in water 
and has a strongly acid reaction. It destroys the micro- 
organisms that cause decay and is used to preserve railroad 
ties and other wooden timbers especially subject to decay. 

Zinc sulfide (ZnS). Very large deposits of zmc sulfide occur 
in southwestern Missouri, constituting the mineral sphalerite. 
It is insoluble in water and, when pure, is white, Lithopone is 
a mixture of the two solids barium sulfate and zinc sulfide, 
made by bringing together barium sulfide and zinc sulfate in 
solution : ^^ + ^^ y B ^ Q ^ + Zng 

It is a valuable white-paint pigment. 

Preservation of wood. With the rapid disappearance of the 
forests the preservation of wood from decay (fungous growths) 
becomes a very important problem. When the wood is to be ex- 
posed merely to atmospheric conditions, it is preserved by paints 







The rubber is obtained 




VI TO 




K Hi JB 


from rubber trees by 




tapping. The rubber, in 




ygff 


the form of an emulsion, 






flows from the trees and 
is collected in buckets, 




,^:':;, / v>W- *.. 


asshowninA. This emul- 
sion resembles milk in 
appearance and is known 




• - 1 
1 ' 1 


as the latex 










JBHH 


The rubber is obtained 




jKbH 


from the emulsion by 




..' JIP^bH 


adding acetic acid, or by 




* ^^W j 1M 


pouring the emulsion, a 




little at a time, over the 




4^ ti ^| 


end of a stick and evap- 




orating the water over a 




fire, as shown in B 








jsiMhB 


"-■' 




.: 'i '^' . i 


yfj 




fe«Br{iL.^BB 










EMMHf 1 ^ 




___ ■ . . ■ ~i 


The raw rubber is next 




«- i~- . i-.. 


mixed (or compounded) 






with sulfur and certain 




, Jl^-^- " 


substances as zinc oxide 
and carbonblack. This 




.«*. 


is done by running the 




■^l" 4* : - ,"/>^ 


materials between large 
revolving cylinders (0), 
as shown in C, until 






thoroughly mixed. Piles 

of the raw rubber are 

shown, marked n 




© ^^B^*" 




— 





Fig. 210, A, B, and C. Some steps in the manufacture of 
automobile casings 



In D cotton fabric is 
being impregnated with 
the compounded rubber 
by running the fabric and 
rubber between large 
metal cylinders. In cord 
tires the cord is wrapped 
about the core and the 
cord as well as the inter- 
stices filled with rubber 



As shown in E, the cas- 
ing is then built up by 
winding the rubberized 
fabric on an iron wheel, 
the surface of which is 
the exact shape of the 
inner part of the desired 
casing. The details can- 
not be shown 



The casing is then placed 
in a mold, as shown in F, 
and heated in large steam 
ovens. In this process the 
rubber is vulcanized and 
toughened. It is then re- 
moved and the finishing 
touches given it 




Fig. 210, D, E, and F. Some steps in the manufacture of 
automobile casings 



350 CHEMISTKY AND ITS USES 

and varnishes. When it must be partly buried in the ground 
(railway ties, fence and telegraph posts), it is treated with sub- 
stances that prevent decay. Those most frequently used are zinc 
chloride, copper sulfate, and creosote from coal tar. 

Cadmium. This element very closely resembles zinc in 
most respects. Some of its alloys are characterized by having 
low melting points. Its compounds are similar in composi- 
tion to the corresponding ones of magnesium and zinc. 

Rubber 

Production. The statement has been made (pp. 109, 349) that 
carbonblack and zinc oxide are largely used in making certain 
rubber goods, and this suggests a brief discussion of rubber itself 
and its application in the making of some important rubber arti- 
cle, such as the casing of an automobile tire. 

Rubber is the product of the rubber tree, which grows wild in 
the upper Amazon region and is cultivated in large numbers in 
tropical regions such as Java, Sumatra, Ceylon, and the Malay 
states. The percentage of rubber obtained from the cultivated 
trees has increased very rapidly as the demand for rubber in- 
creased ; in 1900 only a very small fraction of 1 per cent was 
obtained from this source, while at present nearly 90 per cent of 
the world's output of rubber comes from the rubber plantations. 
The rubber is obtained from the trees by tapping (cutting 
through the bark), as shown in Fig. 210 A. An emulsion runs out 
and is collected in buckets. This fluid consists of water in which 
the rubber particles are suspended, much as the casein and fat are 
suspended in fresh milk; and just as the casein of milk can be 
coagulated by the addition of the acid, so the rubber particles can 
be coagulated by adding acetic acid (p. 295). An older method, 
still in limited use, consists in removing the water by evaporation, 
leaving the rubber. To do this, the emulsion (latex) is poured 
over the end of a stick which is then revolved over the hot smoke 
from a fire (Fig. 210 B). When the water evaporates the process 
is repeated until finally a ball of rubber is built up. 

Composition. Pure rubber is a hydrocarbon whose molecular 
formula is some multiple of the simple formula C 5 H 8 . It can be 



THE MAGNESIUM FAMILY 351 

made in the laboratory, but the process as yet developed is so 
costly that the synthetic rubber cannot compete with the natural 
product. Nevertheless, during the World War the Germans, shut 
off from natural supplies, made a considerable quantity of it and 
even equipped a few automobiles with tires made of synthetic rubber. 

Vulcanizing rubber. The natural rubber is sticky and easily 
affected by the temperature, and as a result is worthless for most 
purposes. In 1839 Goodyear found that if the rubber is heated 
with sulfur these objectionable qualities are overcome. This 
process is known as vulcanization. 

Compounding rubber. For most purposes the pure rubber is 
mixed or compounded with other materials, commonly known 
as fillers. These fillers impart strength or some other desired 
property to the finished product. Carbonblack and zinc oxide are 
the most common of these fillers. The color of a rubber article is 
often due to the filler used. Thus the red color of certain tires 
is due to the use of sulfide of antimony or oxide of iron, while 
the black color of other tires is due to carbonblack. 



EXERCISES 

1. (a) Name the metals so far studied, (b) How many of them will 
displace hydrogen from acids (p. 161) ? 

2. What is the difference in composition between limestone and 
dolomitic limestone? 

3. Write the equation for the reaction (a) between magnesium and 
hydrochloric acid ; (b) between magnesium and dilute sulfuric acid. 

4. What reaction would take place if you added a solution of barium 
chloride to a solution of magnesium sulfate? 

5. (a) Which of the metals so far studied is the heaviest? (b) Which 
has the highest melting point ? 

6. What is the difference in composition («) between Glauber's salt 
and Epsom salt ; (&) between magnesite and Iceland spar ? 

7. Water containing carbonic acid dissolves magnesite just as it 
dissolves limestone. What compound is formed in each case? 

8. Suppose you had a supply of zinc and wished to prepare from it 
the compounds (a) zinc oxide, (b) zinc chloride, (c) zinc sulfide. Give 
the method vou would use. 



352 CHEMISTRY AND ITS USES 

9. Why is it that paint made of zinc oxide is not colored by hydro- 
gen sulfide or other sulfur compounds? 

10. (a) What is an alloy? (b) Mention an alloy of zinc. 

11. Why not use sheets of zinc in place of iron covered with zinc 
(galvanized iron)? 

12. Why is granulated zinc used in preparing hydrogen? 

13. What reaction should you expect to take place if a strip of zinc 
were immersed in a solution of copper sulfate (p. 161) ? 

14. What desirable properties does asbestos possess? 

15. Which ore contains the greater percentage of zinc, sphalerite, or 
zincite ? 

16. What weight of carnallite is necessary for the preparation of 
100 kg. of magnesium? 

17. A zinc plant using the ore franklinite has an output of 5 tons 
of zinc daily. What weight of ore is consumed daily? 

18. Zinc oxide is often made by liberating the zinc from its ores 
and then burning it in air. What weight of sphalerite would be re- 
quired in a plant making 5 tons of zinc white daily? 



CHAPTER XXXIX 

ALUMINIUM 

Introduction. The production of aluminium in quantities 
is a modern accomplishment. In 1883 the production was 
83 pounds, worth about $5 per pound. In 1918 it amounted 
to 225,000,000 pounds, and the 
price was about 30 cents per 
pound. This great advance was 
not due to the discovery of new 
supplies, for the compounds of 
the metal occur hi abundance 
in all our soils, but to improved 
methods for separating the metal 
from its compounds. These 
methods were the outgrowth of 
the work of different investi- 
gators, including the American 
chemist Hall (Fig. 211). In 
1886 Hall, then a student at 
Oberlin College, got the idea 
that it ought to be possible 
to separate aluminium by the 
process of electrolysis, and as 
a result of his investigations the 
method used today and known 
as the Hall process was devised. It is a pleasure to record that 
Hall at his death, in 1914, left several million dollars (a part 
of the fortune gained from his discovery) to Oberlin College, 
where, as a student, he first became interested in the problem. 

353 




Fig. 211. Charles Martin Hall 
(1863-1914) 

The American chemist who developed 
the electrolytic method for the pro- 
duction of aluminium 



354 CHEMISTRY AND ITS USES 

Properties. Aluminium has many desirable properties. It is 
a silver- white metal, only about one third as heavy as iron, 
and yet it is malleable and ductile and fairly hard and strong. 
Moreover, it is an excellent conductor of heat and electricity. 

Occurrence. Next to oxygen and silicon, aluminium is the 
most abundant of all the elements. The free element is not 
found in nature, but its compounds are widely distributed. 
The feldspars, which are the most abundant of all the minerals 
in the earth's crust, are all silicates of aluminium and either 
sodium, potassium, or calcium. Since the soil has been formed 
largely by the disintegration of these rocks, it is rich in the 
silicates of aluminium, chiefly in the form of clay. Some of 
the other forms in which aluminium occurs in nature are the 
following : corundum (A1 2 3 ) ; emery (A1 2 3 colored black 
with oxide of iron); cryolite (Na 3 AlF 6 ); bauxite, a mixture of 
hydrated aluminium oxides (A1 2 3 • H 2 and A1 2 3 • 3 H 2 0) 
together with similar compounds of iron. Bauxite is the ore 
from which aluminium is prepared. In the United States it 
is found chiefly in Arkansas. 

Preparation. The Hall process consists in the electrolysis 
of a solution of aluminium oxide (A1 2 3 ) in melted cryolite. 

An iron box A (Fig. 212) is connected with a powerful elec- 
trical generator in such a way as to serve as the cathode upon 
which the aluminium is deposited. Three or four rows of carbon 
rods B, B dip into the box and serve as the anodes. The box is 
partly filled with cryolite, and the current is turned on, generating 
enough heat to melt the cryolite. Aluminium oxide obtained from 
bauxite is then added, and acts as an electrolyte, being decom- 
posed into aluminium and oxygen. The temperature is maintained 
above the melting point of aluminium, and the liquid metal, being 
heavier than cryolite, collects on the bottom of the vessel, from 
which it is tapped off from time to time through the tap hole C. 

It will be noted that the rather rare mineral bauxite, and 
not the very abundant clay, is used as the source of aluminium. 






ALUMINIUM 



355 



No one has yet succeeded in finding an economical way of 
separating the metal from clay. Here is an opportunity for 
someone to make a discovery equal to that made by Hall. 

Chemical conduct. Neither air nor boiling water has any 
marked effect upon the metal. When heated in oxygen it 
burns with great heat evolution ; hence it is a good reduc- 
ing agent. It dissolves readily in hydrochloric acid and hot 




Fig. 212. Diagram illustrating the manufacture of aluminium 



concentrated sulfuric acid, but nitric acid has little action upon 
it. It is readily dissolved by alkalies and corroded by salt water. 
Uses. The lightness and strength of aluminium, together 
with its inactivity toward air and water, suggest many uses. 
About one third of the output is used in making automobiles. 
Large amounts are also used in the steel industry. Its use 
in the manufacture of kitchen vessels of all kinds is well 
known. Owing to its electrical conductivity it is replacing 
copper to some extent, especially for trolley and power wires. 
In the form of a powder suspended in a suitable liquid it 
makes an efficient, silver-like paint. Many desirable alloys of 
aluminium have been prepared : aluminium bronze (90 per cent 



856 



CHEMISTRY AND ITS USES 



copper and 10 per cent aluminium) has a gold-like color, and 
is strong and permanent in the air. Duraluminum and rnagna- 
lium are also important alloys. 

The World War developed many new uses for aluminium. 
For example, it constituted about one third of the weight of the 
Liberty motor. It was also used in constructing aeroplanes and 
dirigible balloons (Fig. 7). The powdered metal was used in 
certain bombs and explosives. 

Goldschmidt reduction process. Aluminium is frequently 
employed as a powerful reducing agent, many metallic oxides 
which resist reduction by carbon being readily reduced by it. 
The aluminium, in the form of a fine 
powder, is mixed with the metallic 
oxide, together with some substance 
such as fluorite, which melts and thus 
assists by bringing the reacting com- 
pounds in close contact. The mixture 
is ignited, and the aluminium unites 
with the oxygen of the metallic oxide, 
liberating the metal. This collects in 
a fused condition under the melted 
fluorite. 







Fig. 213. Welding car 
rails with thermite 



Thermite welding process. The property 
possessed by aluminium, of reducing ox- 
ides with the liberation of a large amount 
of heat, is turned into practical account in 
the welding of metals, — a method devised by the German chemist 
Goldschmidt. This method may be illustrated by a single exam- 
ple, namely, the welding of car rails, — a process often carried 
out in connection with electric railways to secure good electrical 
connection. The ends of the rails are accurately aligned and thor- 
oughly cleaned. A sand mold A (Fig. 213) is then clamped about 
the ends of the rail, leaving sufficient space so that the metal can 
flow in. The ends of the rails are heated to redness by the flame 
from a gasoline torch directed into the opening in the mold. Just 



ALUMINIUM 



357 



over the opening is placed the crucible B, which contains a mix- 
ture of iron, metallic oxides, and aluminium. When the ends of 
the rails have been heated to redness by the torch, the mixture 
in the crucible is ignited, and after a few seconds the crucible is 
opened at the bottom, and the molten metal resulting from the 
reaction is allowed to flow into the mold. In this way the molten 
metal surrounds the ends of the rails and, as it cools, joins them 
firmly together. A mixture of the metallic oxides and aluminium 
ready for use in weld- 
ing is sold under the 
name of thermite. 

The largest weld 
ever made was that 
of the sternpost of 
the army transport 
Northern Pacific, 
which was broken 
during the World 
War when the ship 
ran aground in a 
fog off Long Island 
(Fig. 214). 

Aluminium Oxide FlG - 214> view °f tlie stern of the Northern 

s ai r\ \ rp-i. -^ Pacific showing the broken post after being welded 

V A1 2 U 3^* lnis sut) ~ by thermite (the break is at the point indicated 
stance occurs in by the arrow) 

several forms in 

nature. The relatively pure crystals are called corundum ; 

emery is a variety colored dark gray or black, usually by iron 

compounds. In transparent crystals, tinted different colors by 

traces of impurities, it forms such precious stones as sapphire, 

ruby, topaz, and oriental amethyst. All these varieties are very 

hard, falling little short of the diamond in this respect. The 

cheaper forms, corundum and emery, are used for cutting and 

grinding purposes. Chemically pure aluminium oxide can be 

made by igniting the hydroxide. The oxide so prepared is 

used in the preparation of aluminium. Some laboratory 




358 CHEMISTKY AND ITS USES 

utensils, such as crucibles and tubes, are made of aluminium 
oxide, which is given the trade name alundum. The same 
material is used for cutting and polishing metals. 

Artificial gems. A number of gems are now prepared in the 
laboratory from molten aluminium oxide. The white sapphires 
so extensively advertised are simply the pure oxide. By fusing 
the aluminium oxide with small percentages of certain other 




Fig. 215. A row of furnaces used in making artificial rubies 
Oxygen is used to obtain high temperatures 

metallic oxides (Fig. 215), different tints or colors are obtained, 
and in this way are prepared such gems as the ruby, the oriental 
amethyst, and the yellow and blue sapphires, which are practically 
identical in composition and properties with the natural stones. 

Aluminium hydroxide (A1(0H) 3 ). The hydroxide can be pre- 
pared by adding ammonium hydroxide to any soluble alumin- 
ium salt, forming a colloidal precipitate which is insoluble 
in water but very hard to filter. 



ALUMINIUM 



359 



Water purification. The value of aluminium hydroxide in the 
purification of water (p. 66) is due largely to its colloidal or gelati- 
nous character when freshly formed by precipitation. After being 
stirred through the water it is allowed to settle slowly, and in so 
doing it carries with it any suspended matter present, including 
microorganisms and coloring materials. Instead of adding alumin- 
ium hydroxide itself to the water, it is more economical and effective 
to produce it by precipitation. This is done by dissolving in the 
water some cheap salt which readily hydrolyzes, such as aluminium 
sulfate : A i 2 ( S 4 ) 3 + 6 H 2 >■ 2 Al(OH) 8 + 3 H 2 S0 4 

There is always sufficient basic material present in the water to 
combine with the sulfuric acid set free, so that no acid is left in the 
water as a result of this treatment. 

Fig. 216 illustrates the use of 
aluminium sulfate in purifying 
water. The cylinder A contains 
impure water. B is a cylinder of 
the same impure water to which 
some aluminium sulfate has 
been added. The aluminium hy- 
droxide formed by hydrolysis 
is slowly settling in the water, 
carrying with it the impurities. 
The appearance of the water 
after settling is shown in C. 

Some cities both purify and 
soften the entire city water- 
supply. In Columbus, Ohio, for 
example, as high as 40,000,000 

gallons is treated daily. The compounds used are (1) aluminium 
sulfate to purify the water, and (2) calcium oxide and sodium car- 
bonate (p. 329) to soften it. These compounds are mixed and 
thoroughly stirred through the water (Fig. 194). A heavy, floccu- 
lent precipitate forms. The water then runs into settling basins 
(Fig. 217; notice the white solid matter in the water as it enters). 
Here it flows slowly forward and backward between walls until 
the precipitate settles. The water is then drawn off, treated with 
chlorine (p. 141), and filtered, 




Fig. 216. Purification of water by- 
aluminium sulfate 



360 



CHEMISTRY AND ITS USES 



Alums. Aluminium sulfate has the property of combining 
with the sulfates of the alkali metals to form compounds 
called alums. Thus, with potassium sulfate the reaction is 
expressed by the equation 

K 2 S0 4 + A1 2 (S0 4 ) 3 + 24 H 2 — ► 2 (KA1(S0 4 ) 2 . 12 H 2 0) 

The sulfates of other tervalent metals can form similar com- 
pounds with the alkali sulfates, and these compounds are also 
called alums, though they contain no aluminium. They all crys- 
tallize in eight-sided crystals and contain 12 molecules of water of 
hydration. The alums most frequently prepared are the following : 

Potassium alum KA1(S0 4 ) 2 • 12 H 2 

Ammonium alum ......... NH 4 A1(S0 4 ) 2 . 12 H 2 

Very large, well-formed crystals of an alum can be prepared by 
suspending a small crystal by a thread in a saturated solution of 

the alum, as shown 
in Fig. 218. The 
small crystal slowly 
grows and assumes 
a very perfect form. 

Hydrolysis of 
salts of aluminium. 
While aluminium 
hydroxide forms 
fairly stable salts 
with strong acids, 
it is such a weak 
base that its salts 
with weak acids 
are readily hydro- 

lyzed (p. 308). Thus, when an aluminium salt and a soluble 

carbonate are brought together in solution, we should expect. 

to have aluminium carbonate precipitated according to the 

equation 2 A1C1 




Fig. 217. Settling basin used in water purification 



3 Na 2 C0 3 



A U C0 «). + 6 NaC1 



ALUMINIUM 



361 



But if it is formed at all, it instantly begins to hydrolyze, the 
products of the hydrolysis being aluminium hydroxide and 
carbonic acid : 



Al 2 (CO s ) 3 + 6H 2 



H 2 CO ! 



>2A1(0H) 3 + 3H 2 C0 3 
* HO + CO, 





■ 


| 




urn 




'KSH;; 






; — ^^^=_ -=_—=!-:! 




<Ji 


^^>-^I 




^ 



Fig. 218. Growing 

a perfect crystal 

of alum 



Aerating agents used in baking. In preparing foods made 
largely from dough, such as bread, biscuits, and cake, it is 
essential that some aerating agent be used 
to render the food light and wholesome. 
The aerating agent used in all cases is 
carbon dioxide. This is generated in the 
dough and, pushing its way through the 
mass, renders it porous and light. The fol- 
lowing methods are used for generating the 
gas in baking : 

1. By alcoholic fermentation. As we have seen, 
this is the method generally used in making 
bread (p. 256). 

2. By the action of sour milk on sodium bicar- 
bonate. The lactic acid present in the sour milk (p. 241) slowly 
acts upon the bicarbonate, liberating carbon dioxide : 

H • C 3 H 5 8 + KaHCO, —+■ NaC 8 H 5 3 + H 2 + C0 2 

3. By the action of an acid salt or alum upon sodium bicarbonate ; 
baking powders. Every baking powder contains three ingredients, 
the name and function of each of which is as follows : (1) sodium 
bicarbonate to furnish the carbon dioxide ; (2) some compound 
which in the presence of water reacts with the bicarbonate and 
slowly liberates carbon dioxide; and (3) starch or flour, which 
keeps the powder dry by absorbing moisture and hence prevents 
deterioration. The compounds used for liberating the carbon 
dioxide from the bicarbonate are either cream of tartar, calcium 
acid phosphate, sodium acid phosphate, or alum. A baking powder 
is generally designated by the name of the compound used to 
liberate the carbon dioxide. Thus, we speak of cream of tartar 
baking powders or alum baking powders. 



362 CHEMISTRY AND ITS USES 

Reactions of baking powders. The reactions that take place 
when water is added to each of the classes of baking powders 
are represented in the following equations : 

Alum (supposing that the alum present is potassium alum) : 

2 KA1(S0 4 ) 2 -f 6 NaHC0 3 ► 

2 Al(OH), + 3 Na 2 S0 4 + K 2 S0 4 4 6 C0 2 
Cream of tartar : 

KHC 4 H 4 6 4 NaHC0 3 >■ KNaC 4 H 4 6 4 H 2 + C0 2 

Phosphate : 

CaH 4 (P0 4 ) 2 + 2 NaHC0 3 ► 

CaHP0 4 4- Na 2 HP0 4 4- 2 H 2 4 2 C0 2 

Dyes and Dyeing 

Properties of dyes. We have already learned (p. 235 ) 
something of the aniline dyes and the history of their dis- 
covery. Because of their very great coloring power these 
dyes have almost entirely taken the place of the dyes ex- 
tracted from plants and trees (Fig. 219). 

The requisites of a good dye are as follows: (1) it must 
have an acceptable ftolor ; (2) it must not injure the fabrics ; 

(3) it must dye fast (in other words, the cloth after having 
been dyed must retain its color when washed with water); 

(4) it must not fade too easily. 

The process of dyeing. This process consists in fixing the 
dye uniformly upon the fabric. The animal fibers, namely, 
wool and silk, are as a rule more readily dyed than cotton, 
which is a vegetable fiber. To dye wool and silk with most 
dyes it is only necessary to steep the fabric in a solution of 
the dye. Cotton fabrics, when treated in this way, will be- 
come colored, but with most dyes the color is not fast. 
Cotton fabrics may be dyed fast in the following way: 

The cloth is first soaked in a solution of an aluminium 
salt (or a similar substance), which readily undergoes 



ALUMINIUM 



363 



hydrolysis. The cloth is then exposed to the action of steam, 
which hydrolyzes the salt, leaving the gelatin-like hydroxide 
thoroughly incorporated in the fiber. The cloth is then steeped 
in the dye, which is absorbed by the aluminium hydroxide 
and is in consequence fastened, or " fixed," upon the fiber. 




Fig. 219. Interior of an indigo factory- 
Indigo, one of our most common dyes, was formerly obtained from a plant grown 
largely in India. Now it is made from coal tar (benzene) at a lower cost 

Aluminium hydroxide and other substances which act in the 
same way are called mordants. The same dye will often give 
different colors with different mordants, as shown upon the 
accompanying colored plate. 

Lakes. The compounds which serve well as mordants may 
be precipitated in solutions containing various dyes, and the 
precipitate will be highly colored, though not always of the 
same color as the dye. Colored precipitates of this kind are 
called lakes and are used as pigments in paints. 



364 CHEMISTRY AND ITS USES 

EXERCISES 

1. State the valences of the metals so far studied. 

2. Enumerate the metals and compounds so far studied that are 
prepared by electrolysis. 

3. Aluminium has been termed the "ideal metal." Give reasons 
for such a claim. 

4. (a) Name the four most abundant elements in the order of 
their abundance (p. 9). (b) Which is the more abundant, iron or 
aluminium ? 

5. Mention two important chemical discoveries made by young men. 

6. Write the equation for the reaction which takes place when 
aluminium dissolves in hydrochloric acid. 

7. Why do the directions for using aluminium cooking utensils 
state that such utensils should not be washed in strong soaps? 

8. Where should you expect factories for the production of alumin- 
ium to be located ? 

9. What substance has largely taken the place of emery for cutting 
and grinding purposes ? 

10. (a) What is the aerating agent used in making bread? (b) Why 
not use baking powder ? 

11. When baking soda is heated, carbon dioxide is evolved (p. 308). 
Why not use it alone as an aerating agent ? 

12. Why does bread have to be set aside to "raise," while biscuits 
and cakes do not ? 

13. Since the passage of the prohibition act cream of tartar baking 
powders are not made in such large quantities. What is the relation 
between prohibition and baking powders ? 

14. What compounds remain in foods as a result of the use of the 
different kinds of baking powders ? 

15. Do you think that a mixture of baking soda and lemon juice 
would act as an aerating agent? 

16. Would baking soda and sweet milk serve as an aerating agent? 

17. Would you classify indigo as a vegetable dye or as a coal- 
tar dye? 

18. An aluminium plant has an output of 2 tons daily. What weight 
of aluminium oxide (A1 2 3 ) would be required daily in the plant? 






ALUMINIUM 365 

19. Which contains the greater percentage of water of hydration: 
potassium alum or borax? 

20. (a) What is the relation between the number of gram-molecular 
weights of baking soda used in a baking powder and the number of 
gram-molecular weights of carbon dioxide generated (see equations, 
p. 362)? (b) Calculate the volume of carbon dioxide generated by 
10 g. of baking soda present in a baking powder. 

21. In what proportion should cream of tartar and baking soda be 
mixed in a baking powder (see equation, p. 362) ? 

22. Ordinary baking powders contain about 12 per cent of starch or 
flour. Calculate the weights of cream of tartar, baking soda, and starch 
required to make 100 g. of cream of tartar baking powder. 

23. (a) Six grams (approximately 2 level teaspoonfuls) of the bak- 
ing powder of the composition referred to in exercise 22 would evolve 
what weight of carbon dioxide? (b) The Federal law requires that a 
baking powder sold on the market must generate 12 per cent of its 
weight of carbon dioxide. Would the baking powder in exercise 22 
meet these requirements? 



CHAPTER XL 



SILICATES AND THEIR COMMERCIAL APPLICATIONS 

Introduction. There are a number of industries in which 
the raw materials used are sand, clay, limestone, and feldspar. 
In the process of manufacture these ingredients form com- 
plex silicates, so that it is convenient to discuss these industries 

under the general term 
of the silicate industries. 
The most important of 
these industries are those 
of glass, cement, and 
clay products, including 
such objects as brick, 
pottery, and chinaware. 
Glass. When cer- 
tain silicates (such as 
those of sodium, potas- 
sium, and calcium) are 
heated together to a 
high temperature with silicon dioxide (sand), the mixture 
slowly fuses to a transparent liquid, which, on cooling, passes 
into the rigid material called glass. In its real nature glass is 
to be regarded as a supercooled liquid rather than as a true 
solid. Instead of starting with the silicates of sodium and 
calcium, it is more convenient and economical to heat sodium 
carbonate (or sulfate) and lime with an excess of clean sand, 
the silicates being formed during the heating: 

Na 2 CO a + Si( ) 2 — y Na 2 Si0 3 + C0 2 
CaO+Si0 2 ^CaSiO s 

366 




Fig. 220. Making a glass object in a mold 




Fig. 221. The manufacture of window glass 

A lump of molten glass is first gathered on the end of a hollow rod (A). By 
using great skill this is then blown (B) into the form of large hollow cylinders. 
These are cut longitudinally (C) and placed in an oven until they soften (D), 
when they are flattened out into plates and cut into the desired shape. The 
comparatively expensive method of blowing the cylinders by mouth is giving 
way to the machine-blown method, the principle of which is as follows : The 
end of a hollow rod is dipped into the molten glass. Air is forced through the 
rod, and at the same time the rod is slowly raised. The molten glass is thus 
formed into cylindrical shape and hardens as it is raised above the molten 
mass. By slowly raising this and forcing in the requisite amount of air to 
prevent the soft glass from collapsing, there are formed cylinders of glass as 
large as two feet in diameter and forty feet in length. These are cut length- 
wise and flattened, as explained above 




Fig. 222. The manufacture of a glass beaker 

The glass-blower first collects the proper amount of molten glass on the end 
of a rod (A). This is then placed in the mold (B) and blown into shape (the 
workman in B is shown removing the beaker from the mold). The glass is 
then cut from the top of the beaker by revolving it in contact with a dia- 
mond (C). Finally, the beaker is heated until slightly soft, when the rim is 
formed, as shown in D 



SILICATES 367 

Molding and blowing glass. The way in which the melted mix- 
ture is handled in the glass factory depends upon the character 
of the object to be made. Many articles, such as bottles, are made 
by blowing the plastic glass into hollow molds of the desired 
shape. The mold is opened, a lump of plastic glass on a hollow 
rod is lowered into it, and the mold is then closed. By blowing 
into the tube the glass is forced into the shape of the mold. The 
mold is then opened (Fig. 220) and the object lifted out. The 
top of the object must be cut off at the proper place and the sharp 
edges rounded off in a flame. Bottles are now more often made 
by machinery, in which the bottle is blown by compressed air. 

Other objects, such as lamp chimneys, glasses, and beakers, 
are revolved while being blown in the mold, and have no ridge 
showing where the mold closes. Window glass has long been made 
by blowing cylinders of glass as explained in Fig. 221. Similar 
cylinders are also made by dipping a tube into melted glass and 
slowly withdrawing it while compressed air is forced into the 
tube. Increasing amounts are also being made by allowing the 
molten glass to flow out onto tables under very definite condi- 
tions. Plate glass is cast into flat slabs, which are then ground 
and polished to perfectly plane surfaces. 

Varieties of glass. By the proper selection of ingredients 
a great variety of glass can be made. Window glass and 
bottle glass are chiefly silicates of sodium and calcium. 
While we think of such glass as insoluble, nevertheless it 
wull not serve for beakers in which strong chemicals are 
heated. For such objects some boric acid is used along with 
the sand, so that the resulting product is a borosilicate of the 
metals present. Pyrex glass, used to withstand shock and 
sudden changes of temperature, is a sodium-aluminium boro- 
silicate containing a large percentage of free silica. Optical 
glass, imported from Germany before the war, but now made 
in the United States, owes its brilliancy to the presence of 
lead silicate. 

Color of glass. The presence of small amounts of certain 
metals impart characteristic colors to glass. Thus, cobalt 



368 



CHEMISTRY AND ITS USES 



compounds give a blue color, while manganese imparts an 
amethyst tinge. Gold produces the finest ruby color, while 
cuprous oxide and selenium give similar ruby tints. The sand 
used in glassmaking must be very pure ; otherwise the iron 
present as an impurity will produce a green color observed in 
bottles made from cheap grades of sand. In some cases the 
color is due to the presence of a colored silicate formed by 

the metal used; in 
others, as in the case 
of gold, the color is 
due to the presence of 
the metal in the form 
of a colloid. 

Aluminium sili- 
cates. When feldspar 
(KAlSi 3 O g ) is decom- 
posed, as explained in 
the chapter on soils 
(p. 338), there are 
formed soluble potas- 
sium compounds, which 
find their way into 
the soil, and aluminium silicate, which is sometimes de- 
posited in beds in nearly pure form known as kaolinite 
(Al 2 Si 2 7 • 2 H 2 0) ; more often it is found mixed with sand 
and other substances, in which form it is known as clay. 
Kaolinite, as well as a clay, when mixed with water, forms a 
plastic mass which can be molded or fashioned into any 
desired shape, and this when baked forms a hard mass ; hence 
their use in making bricks, pottery, and chinaware. Fuller's 
earth is a peculiar form of aluminium silicate which is used as 
a filtering material for decolorizing oils such as cottonseed oil. 
Clay products. The crudest forms of clay products, such 
as porous brick and draintile, have little chemistry involved 




Fig. 223. The manufacture of pottery : mold- 
ing the plastic material into form 



SILICATES 



369 



in their manufacture. Natural clay is molded into the required 
form, dried, and then burned in a kiln, but not to a temper- 
ature at which the materials soften. In this process the nearly 
colorless iron compounds in the clay are converted into colored 
compounds which give the usual red color to these articles. 
In making vitrified 
brick the temperature 
is raised to the point 
at which fusion begins, 
so that the brick is 
partially changed to a 
kind of glass. 

White pottery. This 
term is applied to a va- 
riety of articles, from 
the crudest porcelain 
to the finest chinaware. 
Although the processes 
used in the manufacture 
of the articles differ in 
details, fundamentally 
they are the same and 
may be described under 

three heads ; namely, (1) the preparation of the body of the 
ware, (2) the process -of glazing, and (3) the decoration. 

1. The body of the ware. The materials used consist of an artifi- 
cially compounded clay made from kaolin, plastic clay, and pul- 
verized feldspar. This mixture is plastic and is worked into the 
desired shape by molds or on a potter's wheel (Fig. 223). The 
ware is then dried and burned in a kiln (Fig. 224) until vitrified, 
and in this form is known as bisque. This is usually porous and 
must therefore be glazed to render it nonabsorbent. 

2. The glaze. The glaze is a fusible glass which is melted over 
the surface of the body. The constituents of the glaze are. quartz, 
feldspar, and various metallic oxides, often mixed with a little 




Fig. 224. The manufacture of pottery: stack- 
ing the ware in the kiln for firing 



370 CHEMISTRY AND ITS USES 

boric oxide. These materials are finely ground and mixed with 
water to a paste. Sometimes they are first fused into a glass, 
which is then powdered and made into the paste. The bisque is 
dipped into the glaze paste, dried, and fired until the glaze materials 
melt and flow evenly over the surface. 

3. The decoration. If the article is to be decorated the design 
may be painted upon the body before glazing, or it may be painted 
upon the glaze and the article fired again, the pigments melting 




Fig. 225. A bridge built of reenforced concrete 

into the glaze. In the former case the pigments used are as a rule 
metallic oxides of various colors, while in the latter case they are 
often colored glasses. 

Cement. The term cement as ordinarily used at present is 
applied to those mortars which possess the property of harden- 
ing in water as well as in air. These cements are silicate bodies, 
usually very highly basic in character, and when ground fine 
and mixed with water they undergo complex reactions result- 
ing in the formation of a hard, rocklike mass. A number of 
different classes of cements are known, the most important 
of which, is called Portland cement. 

1 Manufacture of Portland cement. The materials most commonly 

employed are limestone and clay or shale, although any products 

may be used which are similar in composition to the above materials. 

The materials to be used are coarsely ground and then mixed 

together in the proper proportions and finely pulverized. The 



SILICATES 371 

resulting mixture is run into a furnace and burned to a temper- 
ature just short of fusion, at which temperature it vitrifies, form- 
ing a grayish mass known as 'clinker. Finally, the clinker is 
ground to a fine powder. Gypsum is often added in the process ; 
this acts as a negative catalyzer, retarding the hardening, or setting, 
of the cement, 

Importance of cement. The importance of cement in the con- 
struction industries increases each year. It is often used in place 
of mortar in the construction of brick buildings. Mixed with 
crushed stone and sand it forms concrete, which is used in foun- 
dation work for buildings, street-paving, and road-building. It is 
also used in making artificial stone, terra-cotta trimmings for 
buildings, artificial-stone walks and floors, fence posts, and the 
like. It is being used more and more for making articles which 
were formerly made of wood or stone, and ships and the entire 
walls of buildings are sometimes made of cement blocks or of con- 
crete. Iron rods or wire are often embedded in the concrete before 
it sets, to give it greater strength, and this is called reenforced 
concrete (Fig. 225). 

EXERCISES 

1. Sodium silicate alone forms a hard transparent mass. Why not 
use this as glass ? 

2. "What sort of glass would be produced if ordinary impure sand 
were used in place of white sand ? 

3. How can you easily show that window glass has a very pro- 
nounced color? 

4. Why not make glass from sand alone? 

5. What is the difference in composition between feldspar, kaolinite, 
and clay ? 

6. What properties have kaolinite and clay in common that makes 
them adapted for the manufacture of pottery ? 

7. Why is it necessary to glaze chinaware ? 

8. (a) What is the distinction between mortar and cement? 
(6) Could cement be used for laying bricks ? 

9. Give examples of structures you have seen that are made of cement. 
10. Is there any relation in properties between Fuller's earth and 

charcoal ? 



CHAPTER XLI 
THE IRON FAMILY 



Name 


Symbol 


Atomic 
Weight 


Density 


Melting 
Point 


Oxides 


Iron .... 
Cobalt .... 
Nickel. . . . 


Fe 
Co 
Ni 


55.84 

58.97 
58.G8 


7.86 
8.00 
8.90 


1530° 

1480° 
1452° 


FeO, Fe 2 3 

CoO, Co 2 () 3 
NiO, Ni 2 8 



The family. The elements iron, cobalt, and nickel form 
a group in the eighth column of the periodic table. The atomic 
weights of the three are very close together, and their prop- 
erties are very much alike. 

Iron 

Introduction. Next to aluminium, iron is the most abundant 
of the metals. While it is not found free in nature (except in 
meteorites), nevertheless it has long been known and utilized. 
This is due to the fact that it is easy to separate from its 
ores, can be easily molded or hammered into shape, and pos- 
sesses properties that make it useful for many purposes. There 
are many varieties of iron, the properties of which vary because 
of the method of manufacture and the presence of other ele- 
ments, especially carbon. The properties of each of these will 
be described after their manufacture and composition have 
been studied. 

Occurrence. Iron occurs in large deposits as oxides, sulfides, 
and carbonates, and in smaller quantities in a great variety of 
minerals. Indeed, very few rocks or soils are free from small 

372 



THE IRON FAMILY 373 

percentages of iron. It is a constituent of the chlorophyll of 
plants and the haemoglobin of the blood of animals and there- 
fore plays an important part in life processes. Many meteor- 
ites are largely iron in the free state, usually alloyed with a 
little nickel. 

Pure iron. Iron can be prepared in practically pure condi- 
tion by the open-hearth method (p. 379). It is a silvery 
metal which melts at 1530°. It is ductile and malleable and 
is almost as soft as aluminium. It is especially well adapted 
to the manufacture of electromagnets, since it acquires and 
loses magnetic properties more readily than do the ordinary 
varieties of iron. It is also used for purposes where resistance 
to corrosion is desired, for it does not rust rapidly. 

Iron of commerce. Iron differs from most of the other 
metals used in the industries in that the pure metal is only 
prepared to a limited extent and is of limited application, 
while that containing small percentages of other elements 
exhibits a wide variety of properties which make it of the 
greatest value for many different purposes. 

Carbon is always present in amounts which vary from a 
mere trace to about 7 per cent. According to the condition 
of treatment, the carbon may be in the form of graphite 
scattered through the iron, or it may occur dissolved in the 
iron, or as carbides of iron. One of the most important of 
the many carbides of iron is that which has the formula Fe g C 
and is called cementite. Manganese, silicon, and traces of 
phosphorus and sulfur, together with a little oxygen, are 
also present. 

The properties of the iron are greatly modified by the 
percentages of these elements, by their form of combination, 
and by the treatment of the metal during its production. The 
result is that we have many varieties of iron. The most 
common of these varieties recognized in commerce are the 
well-known forms termed east iron, wrought iron, and steel. 



374 



CHEMISTRY AND ITS USES 



Materials used in metallurgy of iron. Four different classes 
of materials are used in the metallurgy of iron : 

1. Iron ore. The ores most frequently employed are the 
following : 



Hematite .„,... Fe 2 () 3 Siderite 
Magnetite F e 3 4 Limonite 



FeC0 3 
2Fe 2 O g 



3 H 2 




Fig. 226. Mining iron ore in Minnesota 



While iron ore is mined in a number of different localities 
in the United States, the great center of production is in the 

neighborhood of Lake 
Superior, the ore being 
chiefly hematite. Large 
amounts are also mined 
near Birmingham, Ala- 
bama. Fig. 226 repre- 
sents one of the large 
mines in Minnesota. 

2. Carbon. Carbon 
in some form is neces- 
sary both as a fuel and 
as a reducing agent. In 
former times wood charcoal was used to supply the carbon, 
but now coke is almost universally used. 

3. Hot air. To maintain the high temperature required for 
the reduction of iron, a very active combustion of fuel is nec- 
essary. This is secured by forcing a strong blast of hot air 
into the lower part of the furnace during the reduction process. 

4. Flux. All the materials which enter the furnace must 
leave it again, either in the form of gases or liquids. The 
iron is drawn off as the liquid metal after its reduction, the 
oxygen with which it was combined escaping as oxides of car- 
bon. To secure the removal of the earthy matter charged into 
the furnace along with the ore, materials are added to the 
charge which will combine with the impurities in the ore, 



THE IRON FAMILY 



375 



forming a liquid. The material added for this purpose is called 
the flux and usually consists of limestone. The liquid pro- 
duced from the flux and the ore is called slag. It is a variety 
of readily fusible glass. 

Cast iron. Ordinarily the first step hi the manufacture of 
any variety of commercial iron is the production of cast iron. 
The ores are mixed with a suitable flux 
and are reduced by heating with coke. 

Blast-furnace process. The reduction is 
carried out in a large tower called a blast 
furnace (Fig. 227). This is usually 80 ft. 
high and 20 ft. in internal diameter 
at its widest part, narrowing somewhat 
toward both the top and the bottom. The 
walls are built of steel and are lined with 
fire brick. The base is provided with a 
number of pipes. 4, called tuyeres, through 
which hot air is forced into the furnace. 
The tuyeres are supplied from a large 
pipe B, which girdles the furnace. At 
the base of the furnace is an opening 
through which the liquid metal can be 
drawn off from time to time. There is 
also a second opening C, somewhat above 
the first, through which the excess of 
slag overflows. The top is closed by a 
movable trap D, called the cone, and 
through this the materials to be used are 
introduced. The gases resulting from the 

combustion of the fuel and the reduction of the ore, together with 
the nitrogen of the air admitted through the tuyeres, escape 
through pipes E. These gases are very hot and contain a sufficient 
percentage of carbon monoxide to render them combustible ; they 
are accordingly utilized for heating the blast of air admitted 
through the tuyeres and as fuel for the engines. 

Charges consisting of coke, ore, and flux in proper proportion 
are at intervals introduced into the furnace through the cone. 




Fig. 227. Vertical section 
of a blast furnace 



376 CHEMISTRY AND ITS USES 

The coke burns fiercely in the hot-air blast, forming carbon dioxide, 
which is at once reduced to carbon monoxide as it passes over the 
highly heated carbon. 

Reduction of the ore begins at the top of the furnace through 
the action of the carbon monoxide. As the ore slowly descends 
the reduction is completed, and the resulting iron melts and 
collects as a liquid in the bottom of the furnace, the lighter 
slag floating above it. After a considerable quantity of iron has 




Fig. 228. Casting pig iron from a blast furnace 

collected, the slag is drawn off through C. The molten iron is 
then run into large ladles or buckets (Fig. 228; A) and used 
directly in the manufacture of steel ; or it is poured from the 
buckets into iron troughs (Fig. 228, B, B) or molds lined with 
lime. These troughs are attached so as to form an endless belt. 
The molten iron in each trough solidifies by the time it reaches 
the farthest point; and as the belt reverses, the solid pieces of 
iron fall into cars placed beneath it. 

In practice a number of furnaces are usually operated together, 
as illustrated in Fig. 229, which shows an exterior view of a modern 
plant for making cast iron. 

Properties of cast iron. The iron produced in the blast fur- 
nace is called cast iron. It varies considerably in composition, 
but always contains over 2 per cent of carbon, variable amounts 



THE IRON FAMILY 



377 



of silicon, and at least traces of phosphorus and sulfur. The 
form in which the carbon is present, whether free or combined, 
also greatly modifies the properties of the iron. In general, 
cast iron is hard and brittle and melts at about 1100°. It 
cannot be welded or forged, but is easily cast in sand molds. 




Fig. 229. A typical plant for the manufacture of cast iron 

It is rigid, but not elastic, and its tensile strength is small. 
It is used for making castings, but chiefly as a starting point 
in the manufacture of other varieties of iron. 

Wrought iron. Wrought iron is made from cast iron by burn- 
ing out most of the carbon, silicon, phosphorus, and sulfur, the 
operation being conducted in what is called a paddling furnace. 

Wrought iron is soft, malleable, and ductile. Its tensile 
strength is greater than that of cast iron, but less than that 
of most steel. Its melting point is much higher than that of 
cast iron. It is no longer produced to the same relative extent 
as in former years, since soft steel can be made at a less cost 
and has almost the same properties. 



378 



CHEMISTRY AND ITS USES 



Steel. Steel, like wrought iron, is made from cast iron by 
burning out a part of the carbon, silicon, phosphorus, and 
sulfur, but the processes used are quite different from that 
employed in the manufacture of wrought iron. Nearly all 
the steel of commerce produced in the United States is made 
by one of two general methods known as the acid Bessemer 
process and the basic open-hearth process. 

Acid Bessemer process. In the acid Bessemer process the 
furnaces used are lined with silica, which, it will be recalled, 
is an acid anhydride. These furnaces 
remove from the cast iron the carbon 
and silicon, but not the phosphorus and 
sulfur. The process is therefore employed 
when the cast iron to be used is low in 
phosphorus and sulfur. 

Details of the Bessemer process. This 
process, invented about 1860, is carried 
out in great egg-shaped crucibles called con- 
verters (Fig. 230), each of which will hold 
as much as 15 tons of steel. The converter 
is built of steel and lined with silica. It is 
mounted on axles, or trunnions, so that it 
can be tipped over on its side for filling 
and emptying. One of the trunnions is hollow, and a pipe con- 
nects it with an air chamber A, which forms a false bottom to 
the converter. The true bottom is perforated so that air can be 
forced in by an air blast admitted through the trunnion and the 
air chamber. 

White-hot liquid cast iron from a blast furnace is run into the 
converter through its open, necklike top B, the converter being 
tipped over to receive it ; the air blast is then turned on and the 
converter turned to a nearly vertical position. The carbon and sili- 
con in the iron are rapidly oxidized (first the silicon and then the 
carbon), the oxidation being attended by a brilliant flame (Fig. 231). 
The heat of the reaction, largely due to the combustion of silicon, 
keeps the iron in a molten condition. The air blast is continued 




Fig. 230. Vertical sec- 
tion showing details of 
a Bessemer converter 






THE IRON FAMILY 



379 



until the character of the flame shows that all the carbon has been 
burned away. The process requires from fifteen to twenty minutes, 
and when it is complete the desired quantity of carbon (generally 
in the form of high carbon iron alloy) is added and allowed to 
mix thoroughly with the 
fluid. The converter is 
then tilted and the steel 
run into molds, and the 
ingots so formed are ham- 
mered or rolled into rails 
or other objects. 

Basic open-hearth 
process. In the basic 
open-hearth process the 
lining of the furnace is 
made of limestone or 
dolomite, both of which 
act as bases. In such 
furnaces the phospho- 
rus and sulfur are both 
removed, as well as 
the silicon and carbon. 
The presence of more 
than traces of phospho- 
rus and sulfur in the 
finished steel renders 
the metal so brittle that 
it is worthless. The open-hearth process therefore possesses 
a great advantage over the' acid Bessemer process in that it 
makes it possible to utilize iron ores (or cast iron obtained from 
them) that contain appreciable quantities of phosphorus and 
sulfur. The operation does not need to be hastened, and steel 
of any desired composition can be produced. An average 
furnace produces about fifty tons of steel in one operation, 
approximately eight hours being required in the process. 




Fig. 231 



The brilliant flame of a Bessemer 
converter 



380 CHEMISTRY AND ITS USES 

In the open-hearth process cast iron is introduced into long, 
low furnaces (Fig 232) lined with limestone or dolomite and 
heated with a gas flame passing above and over the iron. The 
carbon present oxidizes and escapes. The silicon, phosphorus, and 
sulfur are oxidized, and the resulting oxides combine with some 
of the lime lining to form a slag which floats on the surface and 
is easily removed. When the operation is complete the liquid 
steel is drawn off (Fig. 233) into large buckets and then run into 
molds. Most of our steel is made by this process. 

Electrothermal metallurgy of steel. An increasing quan- 
tity of high-grade tool steel is being produced in electrical 
furnaces. The electrical current is used merely to produce 
heat, so that the process is not dependent upon electrolysis. 
This method is almost identical with the open-hearth method, 
save in the way in which the heat is supplied, and produces 
the same kind of steel as does the latter method. 

Properties of steel. Steel contains from a trace up to 2 per 
cent of carbon, less than 0.1 per cent of silicon, and not more 
than traces of phosphorus and sulfur. When desired, a 
product containing as high as 99.85 per cent of iron can be 
produced by the open-hearth method. Such steel is very soft, 
but resists rusting. As the percentage of carbon increases, 
the steel becomes harder and less ductile. Steel can be rolled 
into sheets, cast in molds, and forged into desired shapes. 

The hardening and tempering of steel. When steel contain- 
ing from 0.5 to 1.5 per cent of carbon is heated to a relatively 
high temperature and then cooled suddenly by plunging it 
into cold water or oil, it becomes very hard and brittle. When 
gradually reheated and then allowed to cool slowly, this 
hardened steel becomes softer and less brittle, and this process 
is known as tempering. 

By properly regulating the temperature to which the steel is 
reheated in tempering, it is possible to obtain any condition of 
hardness demanded for a given purpose, as for making springs 




f^hnnnn-gi 




A 



E 



Hot Air J K HotGas 




Fig. 232. Diagram of an open-hearth furnace 

The fuel gas previously heated enters through C, and at I) meets a current 
of hot air entering through B. Vigorous combustion ensues, and the flame 
passes over the cast iron in the furnace, melting it and gradually changing 
it into steel. The hot products of combustion escape through E and F and 
are used in heating the incoming supplies of gas and air. The limestone 
lining' is shown at A, A 





mm Mwm 

JJgL* 

IJKk jjafj^prgHEN §£■■■§ 


Kj^^^ ' "~"7"' 



Fig. 233. Drawing off the molten steel from an open-hearth furnace 
into large buckets, from which it is cast into molds 




Uj 






03 
3 






O o 

«M "f-l 

o s 

CD r-J 



03 



s5sa 

^ 2 « a 

5 ^ ^s 



e« F? 



2 to 



+J o 



? £ J?? fe 



lib! 

""* eg 5 p 



CS3 .w 



CD 



,0 
g n « ^ 

rs * fl s 

5 a .3 a 



2 3 



bn 



O 

"S a .a 

25 I a 



«5 ofl 

o "^ fe- o) 

0) eg -u b 

2 -Q o> 

V, « 2 

B * rt ^ 

§ « -^ 

© "S b~ 

-3 P. o 

^ o » 

el -^ -^ 

g "8 m 

O =*H o 



THE IRON FAMILY 381 

or cutting-tools. Steel assumes different color tints at different 
temperatures, and by these the experienced workman can tell when 
the desired temperature has been reached. Lake gives the follow- 
ing temperatures as suited to the tempering of the tools specified 

220° ...... paper cutters, wood-engraving tools 

210° knife blades, rock drills 

260° hand-plane cutters and cooper's tools 

275° axes, springs 

290° needles, screw drivers 

300° wood saws 

Steel alloys. As we have seen (p. 373), small quantities 
of carbon greatly modify the properties of iron, and equally 
marked effects may be produced by a great many other ele- 
ments. Accordingly, to secure a steel with the requisite prop- 
erties, suitable percentages of these elements are added to the 
steel just before it is run out of the furnace. The elements 
most frequently added are manganese, silicon, nickel, chro- 
mium, tungsten, vanadium, and titanium, and steel contain- 
ing an appreciable percentage of any of these elements is 
called a steel alloy. The element is added in the form of a 
rich alloy of iron, such as ferrochrominm or ferromanganese. 

The approximate composition and uses of some of the 
principal steel alloys are as follows : 

3.5% nickel armor plate 

3.5% nickel and 3.5% chromium .... armor plate and projectiles 

12% manganese burglar-proof safes 

5.0% chromium and from 8 to 24% tungsten high-speed lathe tools 

6.0% chromium and 10% molybdenum . . high-speed lathe tools 

0.1% titanium car rails and steel castings 

0.1% vanadium automobile parts 

12 to 15% silicon retorts for distilling acids 

Compounds of iron. Iron differs from the metals so far 
studied in that it is able to form two series of compounds. 
In the one series the iron is bivalent and forms compounds 
which in formulas and many chemical properties are similar 



382 CHEMISTRY AND ITS USES 

to the corresponding zinc compounds. These are called fer- 
rous compounds. In the other series iron acts as a tervalent 
metal and forms salts similar to those of aluminium. These 
salts are known as ferric compounds. 

Ferrous salts. These salts are obtained by dissolving iron 
in the appropriate acid or, when insoluble, by precipitation. 
The crystallized salts are usually light-green in color. 

Ferrous sulfate (FeSOJ. Ferrous sulfate is the most famil- 
iar ferrous compound. It is usually obtained in the form of 
the hydrate FeS0 4 • 7 H 2 0, called copperas, or green vitriol, 
and is prepared commercially as a by-product in the steel- 
plate mills. Preparatory to galvanizing or tinning (p. 348), 
steel plates are cleaned from rust by immersing them in dilute 
sulfuric acid, and in the process some of the iron dissolves. 
The liquors are concentrated, and the green vitriol separates 
from them. The salt is used in the manufacture of ink and 
iron alum and as a reagent to destroy weeds. 

Ferrous sulfide (FeS). This occurs in nature as the yellowish- 
brown mineral pyrrhotite. In the laboratory it is made by heating 
iron and sulfur together for use in making hydrogen sulfide. 

Iron disulfide (pyrite) (FeS 2 ). This compound occurs in nature in 
the form of brass-yellow cubical crystals often known as fooVs 
gold. When pyrite is burned sulfur dioxide is produced ; hence the 
mineral is used as a source of sulfur dioxide in making sulfuric acid. 

Ferric salts. The crystallized ferric salts are usually yellow 
or violet in color. Heated with water in the absence of free 
acid, they hydrolyze even more readily than the salts of 
aluminium. The most familiar ferric salt is the chloride. 

Ferric chloride (FeCl 3 ). This salt can be obtained most 
conveniently by dissolving iron in hydrochloric acid and then 
passing chlorine into the solution : 

Fe + 2 HC1 ►■ FeCl„ + H 2 

2FeCl 2 + Cl 2 ^2FeCl 3 






THE IROX FAMILY 383 

Ferric hydroxide (Fe(OH) 3 ). When ammonium hydroxide 
is added to a solution of a ferric salt, there is formed a rusty- 
red precipitate which is generally regarded as ferric hydroxide 
(Fe(OH) ). Recent experiments, however, indicate that it is 
the oxide Fe o 3 , associated with a variable amount of water. 

Oxidation of ferrous salts. When a ferrous salt in the presence 
of an acid is oxidized to a ferric salt, it will be noticed that 
the valence of the iron is increased from 2 to 3. This increase in 
valence can often be brought about without the aid of oxygen, as 
is shown in the following equation : 

2 FeCl 2 + Cl 2 >■ 2 FeCl 3 

This is also called an oxidation, although no oxygen takes part 
in the reaction, for the same product is obtained as by the other 
method : 

2 FeCl 2 + 2 HC1 + >■ 2 FeCl 3 + H 2 

In general, when the valence of the metal of a salt is increased, 
the salt is said to be oxidized, whether any oxygen takes part in 
the reaction or not. 

Reduction of ferric salts. Ferric salts may be changed into fer- 
rous salts by the action of nascent hydrogen or other reducing 
agents, as shown in the following equations : 

FeCl 3 + H > FeCl 2 + HC1 

2 FeCl 3 + Zn >■ 2 FeCl 2 +. ZnCl 2 

Although no oxygen is removed in either of these reactions, the 
ferric chloride is said to be reduced ; and, in general, when the 
valence of the metal of a salt is diminished, the salt is said to 
be reduced. 

Sodium ferrocyanide (Na 4 FeC 6 N 6 ) ; potassium ferrocyanide 
(K 4 FeC 6 N 6 ). These are the sodium and the potassium salts of 
the acid H 4 FeC 6 N 6 and are prepared from the by-products 
obtained in the manufacture of coal gas. They are both soluble 
yellow solids, and the potassium salt is often called yelloiv 
prussiate of potash. "When a solution of either is added to a 
solution of a ferric salt such as ferric chloride, a deep-blue 



384 CHEMISTRY AND ITS USES 

precipitate of ferric f errocyanide forms. This is called Prussian 
blue. It is used as a paint pigment and sometimes for bluing 
laundry water. 

Potassium ferricyanide (K 3 FeC 6 N 6 ). By treating a solution of 
potassium ferrocyanide with chlorine water and evaporating the 
solution to crystallization, garnet-red crystals are deposited which 
have the composition K 3 EeC 6 N 6 : 

2 K.FeC fi N fi + CL — >■ 2 KFeCX + 2 KC1 

4 6 6 ' 2 3 6 6' 

This compound is called potassium ferricyanide, or red prussiate 
of potash. 

Blue-printing. When a ferric salt and potassium ferricyanide 
are brought together in solution, no precipitate forms, though the 
solution acquires a yellowish color. On exposure to the sunlight 
the ferric salt undergoes a partial reduction to ferrous salt, and a 
blue precipitate forms. Advantage is taken of these facts in the 
process of blue-printing. A sensitive paper is prepared by soaking 
paper in a solution of potassium ferricyanide and a ferric salt 
(ferric ammonium citrate is generally used), and drying it in a 
dark place. When a black drawing on tracing cloth is placed 
upon such a sensitive paper and the two are exposed to the 
sunlight, the sensitive paper (except where it is protected by the 
black lines) turns a brownish color. It is then thoroughly washed 
with water to remove the soluble salts, during which process the 
portions acted upon by the light turn blue, while the unaffected por- 
tions are left white. A solution of sodium hydroxide can be used as 
an ink for white lettering on a blue-print, since this base decolor- 
izes the blue compound present. 

Cobalt and Nickel 

Occurrence. Cobalt and nickel are almost always found 
together in ores which also contain iron, silver, and copper 
in combination with arsenic and sulfur. The richest deposits 
of cobalt are in Ontario, Canada. Nickel is also a frequent 
impurity of crude copper, and several million pounds of 
nickel sulfate are annually recovered in the United States 
in the refining of copper by electrolysis. 




Iron Chromium Tin Aluminium 

THE USE OF MORDANTS IN DYEING 

The central panel in the figure represents a piece of white cotton cloth. A mor- 
dant of iron, of chromium, of tin, and of aluminium has been applied to the 
cloth in four vertical stripes (leaving the spaces between without a mordant). 
The upper panel shows the same cloth after being dyed with gallein ; the lower 
panel shows a similar piece of cloth after being dyed with alizarin 



THE IRON FAMILY 



385 



Properties and uses. Both these metals are silvery in appear- 
ance and take a high polish. They are somewhat heavier than 
iron and melt at a lower temperature. Their chief use is in 
making alloys. Nickel steel is of the greatest importance and 
is widely used for such purposes as the construction of parts 
of machinery, automobiles, locomotives, armor plate, and pro- 
jectiles. An alloy of cobalt and chromium (stellite) is used 




Fig. 235. The process of electroplating with nickel as carried out in 
a large factory 



for making cutlery and lathe tools. Nickel coinage consists of 
copper and nickel, while German silver contains zinc in addi- 
tion. Nickel is extensively used as a plating upon other 
metals (particularly upon brass) to prevent tarnishing in air, 
and cobalt can be used hi the same way. 

Electroplating with nickel. Nickel-plating is accomplished by 
an electrolytic process. The object to be plated is suspended in a 
solution of a nickel salt and serves as the cathode, while a plate of 
nickel is used as the anode. When the current is passing through 
the electrolyte the nickel is deposited in the form of a silverlike 



W6 CHEMISTRY AND ITS USES 

coating upon the object to be plated, and an equivalent por- 
tion of nickel dissolves from the anode, the composition of the 
electrolyte remaining unchanged. Eig. 235 illustrates the process 
on a large scale, the objects to be plated being suspended from 
the rods A, A. 

Cobalt oxide (CoO). This is the form in which most of the cobalt 
comes into the market. It is a black powder used in making other 
cobalt compounds and in making blue glass and blue decorations 
on china. Sometimes powdered cobalt glass, called smalt, is used 
instead of the oxide and as a pigment. 

Salts of cobalt and nickel. Nearly all the simple salts of cobalt 
and of nickel have formulas similar to those of ferrous salts. The 
most familiar are the following : 

Co(X0 3 ) 2 • 6 H 2 a cherry-red deliquescent salt 

NiS0 4 • 7 H 2 well-formed green crystals 

Inks 

Composition. Inks were known as early as 2500 B. c. and 
were used, in writing. They were black and, like modern inclia 
ink, owed their color to carbon. The composition of ordinary 
black ink may be seen from the following formula, recom- 
mended by the United States government : tannic acid, 
23.4 g. ; gallic acid, 7.7 g. ; ferrous sulfate, 30 g. ; dilate 
hydrochloric acid, 25 g. ; carbolic acid (phenol), 1 g. ; suitable 
blue dye, 2.2 g. ; sufficient water to make 1000 cc. 

Tannic acid and the closely related compound gallic acid 
are obtained from the bark of various trees and especially from 
nutgalls (abnormal growths produced on parts of trees that 
have been stung by certain insects). Ferrous salts do not form 
a color with tannic and gallic acids, but ferric salts produce 
a black tannate and gallate of iron. When the ink is used 
the ferrous salt is oxidized by the air to a ferric salt, and this 
then forms the black iron compound. The dye called for in 
the formula is used to give a temporary color until the per- 
manent color develops. The acid is necessary to keep the iron 



THE IRON FAMILY 387 

in solution. The carbolic acid acts as a preservative to keep 
the ink from molding. Circular No. 9o, issued by the United 
States Bureau of Standards, gives interesting information 
concerning inks. 

Some Special inks. The so-called sympathetic inks were 
used largely during the World War. When such inks are 
used the writing is invisible and remains so until the paper 
is heated or is treated with some substance that will unite 
with the compound composing the ink, to form a colored 
compound. A solution of lead acetate may be used for this 
purpose, the writing being developed by clipping the paper 
into a solution of hydrogen sulfide. A solution of cobalt 
chloride also serves the purpose, the writing being invisible 
until heated gently over a flame. If moistened with water, the 
writing again becomes invisible. Indelible inks are composed 
chiefly of silver nitrate dissolved in dilute aqua ammonia. 
When applied to paper (or fabric) the silver salt is slowly 
reduced to black metallic silver, which imparts the color and 
is not easily removed, since it is insoluble and inactive. For 
the same reason a solution of silver nitrate brought in contact 
with the skin slowly develops a black stain difficult to remove. 

Removal of ink and other stains from textiles. Only some 
general principles can be stated. In the first place, the stain 
should be treated at the earliest moment. Thus, hot water 
will remove fruit stains and many ink stains if used while the 
stain is fresh. If active chemical substances are to be used, 
then a test should first be made upon a small clipping of 
the textile to see that no harm will result. Especially is this 
important if the textile is colored, since the substance used 
may act upon the dye. Fig. 236 shows a convenient way of 
applying chemicals to the stain. 

To be successful one should know the nature of the stain and 
then determine the method of procedure. Some stains may be 
removed by physical methods alone. Thus, fruit stains and coffee 



388 



CHEMISTRY AND ITS USES 



stains may usually be washed out with boiling water, while grease 
spots may be removed by placing the stain over some blotting paper 
and washing it with carbon tetrachloride or low-boiling gasoline. 
The grease is dissolved by the solvent, and the resulting solution 
is absorbed by the paper. If the grease is a solid, such as candle 
grease or paraffin, it may be removed from the fabric by placing 
the stained portion between blotting papers and pressing it with 

a hot iron. The grease 
melts and is absorbed by 
the paper. Turpentine is 
a good solvent for paint 
spots, but must not be 
applied to silk. Many 
substances such as sir- 
ups may be washed out 
with water. 

In many cases it is 
necessary to use chem- 
ical methods. Thus the 
red color produced by 
many acids may be re- 
moved by washing the 
stained portion of the fabric with a little dilute ammonia water. 
Nitric acid acts upon the cloth as well as upon the dye, so that the 
original color cannot be restored. Coffee and fruit stains, if not 
removed by boiling water, may be washed with a mild bleaching 
agent, preferably a solution of sodium hypochlorite, sold by the 
druggist under the name of Javelle ivater. 

Ink stains, if not removed by boiling water, may be treated with 
lemon juice or with a dilute solution of oxalic acid. By this treat- 
ment the ferric salts in the ink are reduced to ferrous salts, which 
can then be washed out with water. Rust stains can be removed 
in a similar way. Some indelible-ink stains may be removed by 
soaking the fabric in a solution of sodium thiosulfate. Silk is so 
sensitive to the action of solvents and reagents that it is generally 
impossible to remove stains from it without injuring the fabric. 

The whole subject of removal of stains from textiles is discussed 
in detail in Farmers' Bulletin 861, published by the United States 
Department of Agriculture. 




Fig. 236. Removing stains on textiles 



THE IRON FAMILY 389 

EXERCISES 

1. Suggest a reason why none of the metals so far studied occur 
free in nature. 

2. The atomic weights of the members of the iron family are 
nearly identical. Are any of their properties correspondingly similar 
(see table at head of chapter) ? 

3. Which is the purest grade of iron? 

4. What are the advantages and the disadvantages of cast iron? 

5. Why is the basic open-hearth process used more than the 
Bessemer ? 

6. How would wrought iron or cast iron do for the construction of 
burglar-proof safes ? 

.7. Why is it necessary to preheat the air blast used in the blast 
furnace but not in the Bessemer converter? 

8. Give the common names of the ordinary hydrates of the sulfates 
of the following metals : sodium, magnesium, iron (ferrous). 

9. How could you tell the difference between fool's gold and real 
gold? 

10. Write the equation for the reaction which takes place when iron 
is dissolved in hydrochloric acid ; in dilute sulfuric acid. 

11. Contrast the relative advantages of iron and aluminium for 
making cooking utensils. 

12. What would be a good test for a ferric salt? 

13. Suggest a reason wh^ some kinds of ink change color (blue to 
black) on drying. 

14. What precautions must one take in removing stains with gasoline ? 

15. (a) Why are objects plated with nickel ? (b) Is there any relation 
between nickel-plated objects and galvanized iron ? 

16. In what important reaction is nickel used as a catalytic agent? 

17. Which contains the greatest percentage of iron : hematite, 
magnetite, or limonite ? 

18. What weight of iron would be required to make 100 kg. of 
copperas ? 

19. What weight of pyrite would be required to make 100 tons of 
50 per cent sulfuric acid ? (Suggestion. 1 gram-molecular weight of pyrite 
makes 2 gram-molecular weights of hydrogen sulfate.) 



CHAPTER XLII 
COPPER, MERCURY, AND SILVER 



Name 


Symbol 


Atomic 
Weight 


Density 


Melting 
Point 


Formulas of 
Oxides 




ous 


ic 


Copper . 
Mercury . . 
Silver . . . 


Cu 
Hg 

Ag 


63.57 
200.60 

107.88 


8.93 
13.56 
10.50 


1083.00° 

- 38.87° 
960.5° 


Cu 2 
Hg a O 

Ag 2 


CuO 
HgO 



The family. Although mercury is not in the same family 
with copper and silver, the three elements resemble each 
other so closely in chemical conduct that it is convenient to 
class them together for study. 

Copper 

Properties and occurrence. Metallic copper has been known 
from the earliest times and was probably the first metal to come 
into considerable use. Its use in the making of instruments 
of war apparently preceded the use of iron. This early use of 
the metal is due to the fact that relatively large quantities of 
it are found in a free state, and in this condition it is easily 
hammered into useful vessels. Its ruddy color also gave it 
value as an ornament. It is a little heavier than iron, but 
melts at a lower temperature (1083°). It is rather soft and 
very malleable and ductile, yet it is tough and strong. As a 
conductor of heat and electricity it is second only to silver ; 
hence its extensive use in the form of wire for conducting 
electrical currents. 

390 



COPPER, MERCURY, AND SILVER 



391 



Copper occurs native in northern Michigan, while its ores 
are found most abundantly in Arizona, Montana, Utah, and 
Michigan. The most valuable ores are the following : 

Chalcopyrite .... CuFeS 2 Cuprite Cu 2 

Chalcocite Cu 2 S Bornite Cu 5 FeS 4 

Nearly all civilized countries produce some copper, but the 
United States produces more than half of the world's supply. 




JFig. 237. Kenning copper by electrolysis 

Chile promises to become the leading copper-producing country 
of the world, large supplies having been opened there recently. 
Metallurgy. The oxide ores are easily reduced with carbon. 
The metallurgy of the sulfide ores is quite complex and beyond 
the scope of an elementary text. It need only be said that the 
copper is first separated in impure form known as blister 
copper, which contains the gold and silver often present 
in copper ores. 



392 CHEMISTRY AND ITS USES 

Refining of copper. If a pure copper is desired, then the blister 
copper must be refined. This is done by electrolysis. A large plate 
of it, serving as an anode, is suspended in a tank, facing a thin 
plate of pure copper which is the cathode. The tank is filled with 
a solution of copper sulfate and sulfuric acid to act as the elec- 
trolyte. An electric current passing through the cell dissolves 
copper from the anode and deposits it upon the cathode in pure 
form, while the impurities, including any gold and silver present, 
collect on the bottom of the tank. Electrolytic copper is one of 
the purest of commercial metals. Fig. 237 shows the pure copper 
cathode raised from an electrolytic cell so that it can be removed 
and the operation repeated. 

Chemical conduct. Since copper is below hydrogen in the 
displacement series, hydrochloric acid and dilute sulfuric acid 
are almost without action upon it ; hot concentrated sulfuric 
acid and nitric acid, however, readily dissolve it (pp. 170, 
193). In moist air it slowly becomes covered with a film 
of the bright-red oxide Cu 2 0, which soon changes to a green 
carbonate. Heated in the air the metal is easily oxidized to 
the black oxide CuO. 

Uses. Copper is extensively used for electrical purposes, 
for roofs and cornices, for sheathing the bottoms of ships, and 
for making alloys. In the following table the composition of 
some of these alloys is indicated: 



Aluminium bronze . . . 90%-98% copper, 2%-10% aluminium 

Brass 63%-73% copper, 27%-37% zinc 

Bronze 70%~95% copper, 1%~25% zinc, 1%-18% tin 

German silver .... 50%-60% copper, 20% zinc, 20%-30% nickel 

Gold coin 10% copper, 90% gold 

Silver coin 10% copper, 90% silver 

Nickel coin . . ' . . . 75% copper, 25% nickel 

Electrotyping. Books are often printed from electrotype plates, 
which are prepared as follows : The face of the type is covered 
with wax, and this is firmly pressed down until a clear impres- 
sion is obtained. The impressed side of the wax is coated with 



COPPER, MERCURY, AND SILVER 393 

graphite, and this is made the cathode in an electrolytic cell con- 
taining a copper salt in solution. The copper is deposited as a 
thin sheet upon the letters in wax and, when detached, is a per- 
fect copy of the type, the under part of the letters being hollow. 
The sheet is strengthened by pouring on the undersurface a suit- 
able amount of commercial lead. The sheet so strengthened is 
then used in printing. 

Two series of copper compounds. Copper, like iron, forms 
two series of compounds : the cuprous compounds, in which it 
is univalent; and the cuprie compounds, in which it is biva- 
lent. The only important cuprous compound is the oxide 
Cu 2 0, which is a bright-red solid sometimes used to impart 
a ruby color to glass. 

Cuprie compounds. Cuprie salts are easily made by dis- 
solving cuprie oxide in acids or, when insoluble, by precipi- 
tation. In crystallized form most of them are blue or green. 
Since they are so much more famjliar than the cuprous salts, 
they are frequently called merely copper salts. 

Cuprie oxide (CuO). This is a black solid obtained by heat- 
ing copper in excess of air or by igniting the hydroxide or 
the nitrate. It is used as an oxidizing agent. 

Cuprie sulfate (CuSOJ. When crystallized from water, 
copper sulfate forms large blue crystals of the hydrate 
CuSO, • 5 HO, called blue vitriol or bluestone. The salt is 

4 2 ' 

a by-product in silver refining and is also made by the 
oxidation of pyrite containing copper sulfide : 
CuS + 2 2 ►■ CuS0 4 

Blue vitriol is used in electrotyping, in copper refining, and in 
the manufacture of insecticides. Like all copper salts it is 
poisonous, especially to lower forms of life. When added, even 
in minute quantities, to water containing green pond scum (algae) 
the plant is quickly killed. The mixture obtained by treating blue 
vitriol with a solution of calcium hydroxide (which precipitates 
copper hydroxide) is called Bordeaux mixture and is used as a 
spray for killing molds and scale on fruit trees. 



394 CHEMISTRY AND ITS USES 

Cupric sulfide (CuS). In the form of a black insoluble pre- 
cipitate cupric sulfide (CuS) is easily prepared by the action 
of hydrogen sulfide upon a solution of a copper salt: 

CuS0 4 + H 2 S — >- CuS + H 2 S0 4 

Mercury 

Properties and occurrence. Mercury stands out prominently 
among the metals for being the only one that is liquid at 
ordinary temperatures. Its use in thermometers and barom- 
eters, under the common name of. quicksilver, has made it 
familiar to everyone. It has long been known and was a 
favorite with the alchemists. Its most common ore is the 
sulfide HgS, known as cinnabar. This occurs in quantities 
in Spain and was mined even as early as Roman times for 
use as a red pigment. The same sulfide, made in the lab- 
oratory by heating mercury and sulfur together, still consti- 
tutes a valuable pigment known as vermilion. Although a 
liquid, mercury is much heavier than iron (density 13.56). 
It boils at 357° and freezes at -38.87°. 

Metallurgy. Mercury is easily separated from cinnabar. 
It is only necessary to roast the ore in air, when the sulfur 
burns, leaving the mercury. Spain leads in the production of 
the metal, followed by California. 

Chemical conduct. When mercury is heated at a low tem- 
perature in the air, it slowly unites with oxygen to form mer- 
curic oxide HgO ; at higher temperatures this decomposes 
into its elements (p. 16). Toward acids mercury conducts 
itself very much like copper. 

Uses. In addition to its use in thermometers and barom- 
eters, mercury is used as a liquid over which to collect gases 
that are soluble in water. It readily combines with metals 
to form alloys called amalgams, and this property is turned 
to account in extracting gold and silver from their ores. 



COPPER, MERCURY, AND SILVER 395 

The common material sold for filling teeth is an amalgam 
of silver and tin. 

Compounds of mercury. Like copper, mercury forms two 
series of compounds : the mercurous compounds, such as 
mercurous chloride (HgCl); and the mercuric compounds, 
represented by mercuric chloride (HgCl 2 ). 

Mercuric oxide (HgO). Mercuric oxide is usually obtained 
as a brick-red substance by carefully heating the nitrate: 

2 Hg(NO a ) 2 — * 2 HgO + 4 N0 2 + 2 

It can also be obtained in a yellow form. 

Mercurous chloride (calomel) (HgCl). Being insoluble, mer- 
curous chloride is precipitated as a white solid when a soluble 
chloride is added to a solution of mercurous nitrate : 

HgN0 3 + NaCl y HgCl + NaN0 3 

It is a much-used medicine. 

Mercuric chloride (corrosive sublimate) (HgCl 2 ). This sub- 
stance is made on a commercial scale by heating a mixture 
of common salt and mercuric sulfate : 

2 NaCl + HgS0 4 ►■ HgCl 2 + Na 2 S0 4 

The mercuric chloride, being readily volatile, vaporizes and 
is condensed again in cool vessels. It resembles mercurous 
chloride in being a white solid, but it is soluble in water. It 
is extremely poisonous, and in dilute solutions is used as an 
antiseptic in dressing wounds. 

Mercuric sulfide (HgS). As cinnabar, this substance forms 
the chief native compound of mercury and occurs in red crys- 
talline masses. By passing hydrogen sulfide into a solution 
of a mercuric salt, mercuric sulfide is precipitated as a Mack 
powder insoluble in water and acids. By other means it can 
be prepared as a brilliant red powder, known as vermilion, 
which is used as a pigment in fine paints. 



396 



CHEMISTRY AND ITS USES 



Silver 

Properties and occurrence. Silver has been known from the 
earliest times and, together with gold, has always ranked as a 
precious metal. The Romans called it argentum and used it for 
ornamental purposes and for coins. Like copper, it occurs 
native as well as in the form of compounds. Most often it 
occurs in combination with sulfur (Ag S) or as a constituent 
of other sulfides. In the United States it is for the most part 
produced in connection with lead. The United States pro- 
duces over one third of the world's output of silver, and the 
American continent about 80 per cent. 

Silver is a heavy, rather soft white metal, very ductile and 
malleable. It is the best conductor of heat and electricity. 
It melts a little lower than copper and, like gold, alloys easily 
with other metals. In the form of a fine powder it is black. 
Chemical conduct and uses. Silver is not acted upon by either 
water or air. Sulfur and its compounds tarnish it, forming 
black silver sulfide (Ag 2 S). Nitric acid is its best solvent, 

forming silver ni- 
trate (AgN0 3 ). 
When its com- 
pounds are re- 
duced in a glass 
vessel the pure 
metal is depos- 
ited as a shining 




Fig. 238. The process of silver plating 



film on the sides of the vessel ; hence its use in making mir- 
rors. Its use for making tableware, ornaments, and coins is 
well known. Silver coins contain 90 per cent of silver and 
10 per cent of copper. 

Electroplating with silver. Since silver is not acted upon by 
water or air and has a pleasing appearance, it is used to coat 
various articles made of cheaper metals. Such articles are said to 



COPPER, MERCURY, AND SILVER 397 

be silver plated, and the process by which this is done is very simi- 
lar to electroplating with nickel (Fig. 235). The object to be plated 
(as, for example, a spoon) is attached to a wire and dipped into a 
solution of a suitable silver salt. Electrical connection is made in 
such a way that the article to be plated is the cathode (Fig. 238), 
while the anode A is made up of one or more plates of silver. 

Compounds of silver. Silver forms only one series of salts, 
which corresponds to the mercurous and the cuprous series. 

Silver nitrate (lunar caustic) (AgN0 3 ). This salt is a white 
solid and is easily prepared by dissolving silver in nitric acid 
and evaporating the resulting solution. When cast into 
sticks it is called lunar caustic, for it has a very corrosive 
action on flesh and is sometimes used in surgery to burn 
away abnormal growths. 

The alchemists designated the metals by the names of the 
heavenly bodies. The moon (luna) was the symbol for silver ; 
hence the name lunar caustic. 

Compounds of silver with the halogens. The chloride, the 
bromide, and the iodide of silver (often termed collectively 
the silver halides) are insoluble in water and in acids and 
therefore are precipitated by bringing together a soluble 
halogen salt with silver nitrate: 

AgX0 3 + KC1 — y AgCl + KNO a 

Tliey are remarkable for the fact that they are very sensitive to 
the action of light, undergoing a change of color and chemical 
composition when exposed to sunlight, especially if in contact 
ivith organic matter, such as gelatin. It is upon this property 
of the silver halides that the art, of photography is based. 

Photography. From a chemical standpoint the processes of pho- 
tography may be described under two heads : (1) the preparation 
of the negative ; (2) the preparation of the print. 

1. Preparation of the negative. The plate used in the preparation 
of the negative is made by spreading a thin layer of gelatin, 
in which colloidal silver bromide is suspended (silver iodide is 



398 CHEMISTRY AND ITS USES 

sometimes added also), over a glass plate or more often a nitro- 
cellulose film (Fig. 150) and allowing it to dry. When the plate 
so prepared is placed in a camera and the image of some object is 
focused upon it, the silver salt undergoes a change (not well 
understood) which is proportional at each point to the intensity 
of the light falling upon it. In this way an image of the object 
photographed is produced upon the plate. This image, however, 
is invisible and is therefore called latent. It can be made visible 
by the process of developing. 

To develop the image the exposed plate is immersed in a solu- 
tion of some reducing agent called the developer. While the de- 
veloper will in time reduce all the silver salt, it acts much more 
rapidly upon that which has been exposed to the light. The action 
is therefore continued only long enough to bring out the image. 
The reduced silver is deposited in the form of a black film which 
adheres closely to the plate. 

The unaffected silver salt is now removed from the plate by 
immersing it in a solution of sodium thiosulfate {hypo) (p. 310). 
The plate is then washed with water and dried. The plate so 
prepared is called the negative (Fig. 241) because it is a picture 
of the object photographed, with the lights and positions exactly 
reversed. 

2. Preparation of the print. The print is made on paper which is 
prepared in much the same way as the negative plate. The nega- 
tive is placed upon this paper and exposed to the light in such a 
way that the light must pass through the negative before striking 
the paper. If the paper is coated with silver chloride a visible 
image is produced, in which case a developer is not needed. 
Proofs are made in this way. In order to make them permanent 
the unchanged silver chloride must be dissolved off with sodium 
thiosulfate. The print is then toned by dipping it into a solution 
of gold or platinum salts, in which process the silver on the print 
passes into solution, while the gold or platinum takes its place. 
These metals give a characteristic color or tone to the print, the 
gold making it reddish brown, while the platinum gives it a steel- 
gray tone. Since the darkest places on the negative cut off the 
most light, it is evident that the lights of the print (Fig. 242) 
will be the reverse of those of the negative and will therefore 
correspond to those of the object photographed. 







Fig. 239. Dissolving silver in nitric acid preparatory to making the 
silver halides used in photographic films 

One company uses 6000 pounds of silver weekly for this purpose 




Fig. 240. Nitrating cotton for the manufacture of photographic films 

The cotton is led down through the inclined tubes into the kettles, which 
contain a mixture of nitric and sulfuric acids 




Ihe negative plate 



The photographic plate or film, after exposure in the camera, is immersed 
in a solution of a reducing agent (the developer) which reduces the silver 
salt that has been acted upon by the light to black metallic silver. The un- 
affected silver salt is then dissolved, leaving a picture of the object photo- 
graphed on the plate, but with the lights and positions exactly reversed. 
The plate so prepared is called the negative 




Fig. 242. The positive print 

The print is made from the negative. The photographic paper on which 
the print is to be made is placed under the negative and exposed to the 
light. The paper is then treated with a developer, and the unaffected silver 
salts are dissolved off as in the preparation of the negative. Since the 
darkest places on the negative cut off the most light, the lights of the 
print will be the reverse of those of the negative ; in other words, they 
will correspond to those of the object photographed 



COPPER, MERCURY, AND SILVER 399 

EXERCISES 

1. Point out some respects in which copper and mercury resemble 
each other. 

2. Note the position of copper, mercury, and silver in the displace- 
ment series. What facts concerning these metals could you predict from 
their position in the series ? 

3. (a) Which of the metals studied are found in a free state in 
nature ? (b) What is their position in the displacement series with 
reference to hydrogen ? 

4. (a) Which of the metals so far studied is the heaviest? the 
lightest ? (/>) Which has the highest melting point ? the lowest melting 
point? (Consult table in Appendix.) 

5. What metal other than copper is used in making wires for 
conducting electricity ? 

6. Give the composition of a nickel coin ; a silver coin ; a gold coin. 

7. What is the distinction in composition between brass, bronze, 
and German silver? 

8. What is the distinction in composition between blue vitriol 
and copperas? 

9. What are the properties of mercury that make it well adapted 
for use in thermometers and barometers ? 

10. What are the extremes of temperature that can be registered by 
an ordinary thermometer? 

11. Some thermometers have the space above the mercury filled with 
nitrogen. Can you suggest any reason for this ? 

12. Give the composition of each of the following : calomel, corrosive 
sublimate, lunar caustic, vermilion. 

13. Suppose you got mercury in contact with a gold ring. What 
would be the result ? 

14. A solution of copper sulfate reacts acid to litmus. Explain. 

15. What will happen if you put your knife blade into a solution 
of copper sulphate ? 

16. The alchemists claimed that by putting iron into copper sulfate 
solution the iron was changed into copper. Point out the fallacy. 

17. Why do silver spoons blacken when they come in contact with 
certain foods ? 



400 CHEMISTRY AND ITS USES 

18. The shutter, on a camera did not open, so that the film was not 
exposed. Not knowing this the photographer proceeded with the film 
as usual. What results would he obtain? 

19. A photographic film was accidentally exposed to the light before 
it was placed in the camera. What results would be obtained when the 
plate was developed and treated in the usual way ? 

20. What are the steps in making silver bromide from the metallic 
silver? 

21. What is the disadvantage in using nitrocellulose for making 
photographic films ? 

22. Suppose you wish to electroplate a metal tablet with copper, how 
would you proceed? 

23. How could you easily distinguish between mercuric oxide and 
cupric oxide? 

24. (a) Explain the chemistry involved in toning a print (p. 161). 
(b) Could a print be toned by placing it in a solution of copper salt? 

25. Suppose you were given the following substances : sulfuric acid, 
hydrochloric acid, nitric acid, marble, copper, and zinc. What gases 
could you prepare from these substances? 

26. Suppose you wish to prepare calomel and corrosive sublimate on 
a large scale, what raw materials would you require ? 

27. Suppose you were using chalcopyrite as a source of copper and 
wished to obtain 100 tons of the metal, what weight of the ore would 
be required ? 

28. What is the percentage of mercury in cinnabar? 

29. A silver dollar weighs approximately 26.5 g. What weight of 
silver nitrate could be made from such a coin ? 

30. It has been stated that a single photographic plant uses 6000 lb. 
of silver weekly for photography. What weight of silver bromide would 
this quantity of silver make ? 



CHAPTER XLIII 
TIN AND LEAD 



Name 


Symbol 


Atomic 
Weight 


Density 


Melting 
Point 


Common Oxides 


Tin 

Lead .... 


Sn 
Pb 


118.70 
207.20 


7.30 
11.37 


231.9° 
327.4° 


SnO m Sn0 2 
Pbo'Pb 3 4 Pb0 2 



Tin 

Occurrence. Tin has long been used and is familiar to all of 
us. The ancients confused it with lead, but about the begin- 
ning of the Christian era the distinction between the two came 
to be recognized. The element occurs in nature chiefly in the 
form of the oxide Sn0 2 , which is known as cassiterite or tin 
stone. This is found principally in Malaya, Bolivia, Banka, 
and China and is the ore from which our supply of the metal 
comes. Practically none is produced in the United States. 

Properties. Pure tin, called block tin, is a soft white metal 
with a silverlike appearance and luster ; it melts readily and 
is somewhat lighter than copper. It is malleable and can be 
rolled out into very thin sheets, forming tin foil ; most tin 
foil, however, contains a considerable percentage of lead. 

Metallurgy. The metallurgy of tin is very simple. It is 
only necessary to heat the ore with carbon, which removes 
the oxygen : 



Sn0 2 + C 



>Sn + CO ( 



Chemical conduct. Under ordinary conditions tin is un- 
changed by air or moisture, but at a high temperature it burns, 

401 






402 CHEMISTRY AND ITS USES 



forming the oxide Sn0 2 . Dilute acids have little effect upon 
it, but concentrated acids attack it readily. Concentrated 
hydrochloric acid changes it into the chloride SnCl 2 . 

Uses of tin. A great deal of tin is used in the making of 
tin plate. The process consists in dipping thin sheets of iron 
into the melted tin and is quite similar to that of galvanizing 
iron (p. 348). Owing to its resistance to the action of air 
and weak acids, tin plate is used in many ways, such as in 
roofing and in the manufacture of tin cans, cooking vessels, 
and similar articles. Small pipes of block tin are used instead 
of lead for conveying pure water or liquids containing dilute 
acids, such as soda water. Many useful alloys contain tin. 
Pewter and soft solder are alloys of tin and lead. 

Soldering and brazing. The use of solder in joining two metal 
surfaces depends upon (1) the low melting point of the solder, 
and (2) the fact that it flows over clean metal surfaces and sticks 
to them on cooling. To secure clean surfaces free from oxide a 
suitable flux must be used which will either dissolve the oxide as 
fast as it forms or will reduce it again to metal. The usual fluxes 
are zinc chloride, ammonium chloride, rosin, and stearin. In braz- 
ing or hard soldering the process is essentially the same except 
that a low-melting brass is used instead of solder, and borax is 
used as a flux (p. 291). 

Compounds of tin. Tin forms two series of metallic com- 
pounds : the stannous, in which the tin is bivalent, as is illus- 
trated in the compounds SnO, SnS, SnCl ; and the stannic, 
in which it is quadrivalent, as shown in the compounds Sn0 2 , 
SnS 2 . 

Chlorides of tin. Stannous chloride is prepared by dissolv- 
ing tin in concentrated hydrochloric acid and evaporating 
the solution to crystallization. If metallic tin is heated in a 
current of dry chlorine, anhydrous stannic chloride (SnCl 4 ) is 
obtained as a heavy, colorless liquid which fumes strongly 
on exposure to air. A great deal of tin in the form of stannic 



TIN AND LEAD 403 

chloride is recovered from scrap tin by the action of chlorine. 
From solution in water a solid hydrate is obtained of the 
formula SnCl 4 ■ 5 H 2 0. 

The chlorides of tin are much used as mordants in dyeing 
processes, in calico printing, and to give weight to silk. 

In the World War hand grenades containing an explosive 
charge and stannic chloride were used to clear dugouts of enemy 
troops. The explosion of the grenades filled the dugouts with a 
fine mist of the chloride so that one could not breathe the air 
without painful suffocation. 

Lead 

Properties. Articles made of lead have been found in very 
old Egyptian ruins, and there is no doubt but that it was used 
from early times. The Romans called it plumbum and used it 
for water pipes, as we do today. It is a heavy metal which 
has a bright luster on a freshly cut surface but soon tarnishes 
on exposure to air. It is soft, easily melted, and has little 
strength. 

Occurrence. The United States produces about one third 
of the world's supply of lead, the chief lead-producing states 
being Missouri, Idaho, Utah, and Colorado. The chief ore is 
galenite (PbS) (Fig. 243), and it usually contains some silver. 

The metal is separated by a rather complicated process of roast- 
ing and reduction, and the crude lead is alloyed with any silver 
present in the ore and also with smaller percentages of other 
metals with which galena is associated. It is partially purified 
by melting it in a furnace with free access of air. Some of the 
impurities are thus changed into oxides, which float on the surface 
of the molten lead and so can be removed. The silver, however, 
still remains alloyed with the lead, from which it is separated by 
suitable processes. 

Chemical conduct. When lead is heated in the air it forms a 
yellow oxide (PbO). With the exception of hydrochloric and 
sulfuric acids (which form insoluble compounds), most acids, 



404 CHEMISTRY AND ITS USES 



even very weak ones, act upon the metal, forming soluble 
lead salts. Hot concentrated hydrochloric and sulfuric acids 
also attack it to a slight extent. 

Uses. Lead finds many important applications in the in- 
dustries, chiefly in the manufacture of storage batteries, in 
linings for sulfuric acid plants, in alloys of various kinds 
(such as shot, antifriction metals, type metal, and pewter), 







Fig. 243. A crystal of galenite embedded in calcite 

and in water pipes for plumbing. Since lead dissolves to 
some extent in pure water, it should not be used for pipes 
that are to carry rain water. About one third of the annual pro- 
duction of lead is used in making paint and is permanently lost. 

Compounds of lead. In nearly all its compounds lead is 
bivalent, but in a few of its compounds it has a valence of 
four. All its compounds are poisonous to some extent. 

Lead oxides. Lead forms a number of oxides, the most 
important of which are the following: 

1. Litharge (PbO). This oxide forms when lead is oxidized 
at a rather low temperature and is obtained as a by-product 



TIN AND LEAD 



405 




in silver refining. In color it ranges from yellow and light 
brown to red, depending on its mode of production. It has 
a number of commercial uses. 

2. Red lead, or minium (P& 3 4 ). Min- 
ium is prepared by heating lead (or 
litharge) to a high temperature in 
contact with a current of air. It is a 
heavy powder of a beautiful red color 
and is much used as a pigment for 
painting structural iron. 

Lead sulfide (PbS). In nature this 
compound occurs in a highly crystalline 
form called galenite (Fig. 243), the crys- 
tals having much the same color and 
luster as pure lead. It is readily pre- 
pared in the laboratory as a black pre- 
cipitate, by the action of hydrogen sulfide upon soluble lead 

Pb(N0 3 ) 2 + H 2 S — y PbS + 2 HN0 8 

Lead carbonate. The normal carbonate of lead (PbC0 3 ) 
is found to some extent in nature and can be prepared in 

the laboratory ; a basic 

carbonate, however, can 
be obtained much more 
easily. This is a com- 
pound of lead carbon- 
ate and lead hydroxide. 
It has the formula 



Fig. 244. A crock filled 

with thin lead plates for 

making white lead 



(PbCO s ) .Pb(OH), 




Fig. 245. Lead buckles before and after 
exposure to acetic acid and carbon dioxide 



and is commonly called 
white had. It is pre- 
pared on a very large scale as a white pigment and as a body 
for paints which are to be colored with other substances. 



406 



CHEMISTRY AND ITS USES 



Manufacture of white lead. White lead can be prepared by 
number of processes, but no other seems to produce a product oi 
as desirable physical properties as the old Dutch process, which 
has been used for centuries, though with many improvements. In 
this process the lead is cast into perforated plates called buckles, 
which are placed loosely upon each other in a crock of the shape 




Fig. 246. Harvesting flax 

Its seeds when pressed give us linseed oil, while the residue is a valuable 
cattle food. Linen is made from the steins of the plant 



shown in Fig. 244, the ledge B formed by the constriction of the 
crock supporting the plates. Under them, in A, is poured a suitable 
quantity of dilute acetic acid, and the crocks so charged are placed 
in banks and covered with stable manure or spent tanbark. The 
heat of fermentation in the latter warms the acid, the fumes of 
which attack the lead, forming lead acetate. The carbon dioxide 
from the fermentation enters into reaction with the acetate and 
produces the basic carbonate, regenerating acetic acid, which acts 
again upon the lead. The process continues until the buckles are 
almost completely converted into the desired compound. Fig. 245 
shows a buckle before and after corrosion. 



TIN AND LEAD 



407 



OTHER IMPORTANT COMPOUNDS OF LEAD 

Lead nitrate (Pb(N0 3 ) 2 ) : white soluble crystals 

Lead chloride (PbCl 2 ) : white needles, very sparingly soluble 

Lead sulfate (PbS0 4 ) : an insoluble white crystalline powder 

Lead acetate (Pb(C 2 H 3 2 ) 2 • 3 H 2 0) : a soluble white salt called sugar 

of lead 
Lead chromate (PbCr0 4 ) : used as a pigment in paint (chrome yellow} 



Paints 

Composition. A paint consists of three essential ingredi- 
ents : the vehicle, the body, and the pigment. 

1. The vehicle, or liquid medium. This must be an oil which will 
dry rapidly and harden in drying to a more or less flexible, hornlike 
body. These changes in the 
oil are due to oxidation by 
the air. A number of dif- 
ferent oils will serve this 
purpose, but linseed oil 
(Fig. 246) has long been 
used as the standard dry- 
ing oil, since it can be pro- 
duced in quantity and at 
moderate cost. It is cus- 
tomary to add to it a dryer, 
made by boiling some of 
the oil with oxides of man- 
ganese, lead, or cobalt. The 
oxides enter into combi- 
nation with the oil and 
assist catalytically in its 
oxidation. 

2. The body. The body 
of the paint must be some 
solid material, suspended 

in the oil, which will give a smooth and waxy surface as the paint 
dries and will have good covering power. While white lead meets 
these requirements, it is moderately expensive, and it also blackens 



v 1 •"' 


■' 






ss, II » ■ 



ElG. 



247. Factory appliances used in 
mixing and grinding paint 



408 



CHEMISTRY AND ITS USES 



,1. J : ,1 . * 


.«■ 1 •' 1 . 

IV 'Hi 


Sk 





Fig. 248. Making incisions in pine trees 

From these incisions exudes a resinous mass 
which is the raw material from which ordi- 
nary rosin and oil of turpentine are prepared 



cases they are natural 
products. Sometimes 
they are prepared by 
precipitating an amor- 
phous body (usually a 
colloid) in the presence 
of an organic dye, the 
dye being adsorbed by 
the precipitate and giv- 
ing it a color. Such 
pigments can be pre- 
pared in endless variety 
of colors and are called 
lakes. They are usually 
not so permanent as min- 
eral pigments but serve 
the purpose very well. 



when exposed to hydrogen 
sulfide, which is likely to be 
present in the air in cities. 
Other bodies are now fre- 
quently combined with the 
lead or replace it altogether, 
among them being zinc oxide, 
barium sulfate, and a product 
called lithopone (p. 349). For 
some purposes these mate- 
rials are a real advantage, 
and they are not to be re- 
garded as adulterants unless 
sold as white lead. 

3. The pigment, or coloring 
matter. In the case of white 
paints the body serves also 
as the coloring matter. For 
other colors a specific pigment 
must be added. Frequently 
these pigments are metallic 
oxides or salts and in most 




Fig. 249. The resinous mass (1) which exudes 
from pine trees (Fig. 248), when distilled, 
yields oil of turpentine (3), sometimes known 
as spirits of turpentine, and leaves a residue (2) 
which is ordinary rosin, used in making 
cheap soaps and for many other purposes 



TIN AND LEAD 



409 



Fig. 247 represents the method of manufacture of paint. The 
body, together with a little oil, enters at A and is ground in suc- 
cession in B, C, D, and E, during which process the requisite 
amounts of oil, dryer, and pigment are added. 

Varnish. Varnish is a liquid which, on being applied to a sur- 
face and left to stand, forms a closely adhering and generally 







4 

saJBt 


:: j^fc ■ P*P- 




%$&*^ 



Fig. 250. Making varnish 

The resin is placed in large iron kettles set on wheels, and melted over a fire. The 

kettles are then drawn away from the fire, and the resin is dissolved in linseed oil 

together with oil of turpentine or a similar liquid, such as henziue 



transparent film. Tarnishes are made by dissolving in appropriate 
solvents the resins (or gums) obtained from certain trees. There are 
two chief kinds : (1) In the one the resin is dissolved in linseed 
oil (Fig. 250) thinned with spirits of turpentine or a similar liquid, 
such as benzine. On exposure to the air the turpentine evapo- 
rates and the linseed oil oxidizes and dries as in paints. (2) In 
the other class, known as spirit varnishes, the resin is dissolved 
in some volatile liquid, as spirits of turpentine or alcohol. On 
exposure to air the solvent quickly evaporates, leaving the resin. 



410 CHEMISTRY AND ITS USES 

Spirits of turpentine and rosin (which is not a true resin, but is 
sometimes used in varnish) is obtained from pine trees, especially 
from a species that grows in the Southern states (Figs. 248, 249). 

EXERCISES 

1. How could you distinguish between tin and lead? 

2. (a) What commercial products consist of some one metal coated 
with a layer of a second metal ? (b) What is the object of using the 
two metals in each case? 

3. AVhat salts other than those of tin are nsed as mordants? 

4. Which of the compounds of tin is a liquid? 

5. How could you distinguish between lead oxide (PbO) and mer- 
curic oxide (HgO) ? 

6. How can you remove paint from clothing? 

7. Mention different substances studied that are obtained from trees 
of various kinds. 

8. Which of the lead compounds studied have a color? 

9. Which of the metals studied are not produced to any extent in 
the United States? 

10. What properties have spirits of turpentine, alcohol, and benzine 
in common that ^dapt them for use in making varnish ? 

11. Are there any objections to using benzene as a solvent in making 
varnish ? 

12. Linseed oil is more expensive than cottonseed oil. Why not use 
the latter in making paint? 

13. Painters using white lead often suffer from lead poisoning. 
Magnesium sulfate is said to be a good antidote for lead poisoning. 
Explain its action. (Suggestion. A compound must be soluble in order 
to be absorbed in the system.) 

14. Mention an alloy of tin ; of lead. 

15. For what purposes is rosin used in addition to making varnish ? 

16. What weight of tin will one ton of cassiterite give on reduction ? 

17. Suppose you wish to prepare 100 kg. of stannic chloride for use 
in silk manufacture, what weight of tin will be required? 

18. What is the per cent of lead in galenite? 






CHAPTER XLIV 
MANGANESE AND CHROMIUM 



Name 


Symbol 


Atomic 
Weight 


Density 


Melting 
Point 


Formulas of Acids 


Manganese . 
Chromium 


Mn 
Cr 


54.93 
52.00 


8.01 

7.30 


1230° 
1615° 


H 2 Mn0 4 and HMn0 4 
H 2 Cr0 4 and H 2 O 2 7 



General. Manganese and chromium, although in different 
periodic families, may be discussed together. The metals them- 
selves are important constituents of some of the steel alloys. 



Manganese 

Occurrence and properties. It ^vill be recalled that in the 
preparation of both oxygen and chlorine the oxide MnO., 
was used. This compound occurs in nature as the mineral 
pyrolusite and is the ore from which manganese and its com- 
pounds are prepared. The largest deposits are in Brazil and 
India, but during the World War, when it was difficult to 
secure adequate supplies from these countries, considerable 
manganese ore was mined in the United States. 

The metal is hard and brittle and looks somewhat like iron. 
Its density and melting point are not far from those of iron. 
It liberates hydrogen from dilute acids and from water. 

Metallurgy. The pure metal is best prepared by the Gold- 
schmidt process (p. 356). By reducing a mixture of oxides 
of iron and manganese there are obtained alloys of the two 
metals, known as sjner/el iron and ferromanganese. 

411 



412 CHEMISTRY AND ITS USES 






Compounds of manganese. Manganese not only forms salts 
like the other metals but also forms unstable acids, some of 
whose salts are of great importance. 

Manganese salts. Manganese resembles iron in that it acts 
both as a bivalent and a tervalent metal. The manganous salts, 
in which the metal is bivalent, are the only important ones. 
These have formulas similar to the ferrous salts and resemble 
them in many of their chemical properties. Most of them are 
pink in color. 

Permanganic acid (HMn0 4 ) ; potassium permanganate 
(KMnOJ. Potassium permanganate is a salt of the unstable 
permanganic acid (HMn0 4 ). It forms deep-purple crystals 
readily soluble in water. It is a powerful oxidizing agent 
and is used for this purpose both in the laboratory and as a 
disinfectant and an antiseptic. 

Oxidizing properties of potassium permanganate. Potassium per- 
manganate is remarkable for its strong oxidizing properties, espe- 
cially in the presence of an acid. When sulfuric acid is present 
the reaction takes place in such a way that both the potassium and 
the 'manganese are changed into sulfates, with the liberation of 
oxygen, as shown in the equation 

2 KMn0 4 + 3 H 2 S0 4 >- K 2 S0 4 + 2 MnS0 4 + 3 H 2 +50 

Under ordinary conditions, however, the reaction does not take 
place unless a third substance is present which is capable of oxida- 
tion. The oxygen is not given off in the free state, but is used up 
in effecting oxidation. 

Chromium 

General. Chromium is a hard metal about as heavy as 
iron. The only important ore is chromite or chrome iron ore, 
which has the composition FeCr 2 4 . This ore is found chiefly 
in Rhodesia, New Caledonia, and Greece. A small amount 
occurs in California. The metallurgy of the metal is similar 
to that of manganese. 



MANGANESE AND CHROMIUM 413 

Compounds of chromium. Like manganese, chromium acts, 
both as a base-forming and an acid-forming element. Nearly 
all its compounds are colored ; hence the name chromium, which 
is derived from a word meaning " color." • 

Chromic hydroxide (Cr(OH) 3 ); chromic salts. Chromic 
hydroxide is a greenish compound. It is insoluble and can 
therefore be prepared by precipitation. It dissolves in acids, 
forming the corresponding chromic salts. These are green or 
violet in color and have formulas similar to the ferric salts. 
They are used as mordants (p. 363). 

Chromates. The chromates are salts of the unstable chromic 
acid (H 2 Cr0 4 ) and as a rule are yellow in color. Most of 
the chromates are insoluble and can be prepared from the 
soluble potassium salt by precipitation. In the case of lead 
chromate the equation is as follows: 

Pb(N0 3 ) 2 + K 2 Cr0 4 >- PbCr0 4 + 2 KN0 3 

Lead chromate (chrome yellow') and barium chromate are used 
as yellow pigments. 

Potassium dichromate (K 2 Cr 2 7 ). When potassium chromate 
is treated with sulfuric acid, the potassium salt of dichromic 
acid (H 2 Cr 2 7 ) is formed: 

2 K 2 Cr0 4 + H 2 S0 4 — *- K,Cr,0, + K 2 S0 4 + H 2 

This is the best-known dichromate and the most familiar 
chromium compound. It forms large brilliant crystals of a 
red color and is rather sparingly soluble in water. Potas- 
sium dichromate, as well as the corresponding sodium salt 
(Na 2 Cr 2 ? ), finds use in many industries as an oxidizing 
agent, especially in the preparation of organic substances and 
in the construction of several kinds of electrical batteries. 

Oxidizing action of potassium dichromate. When a dilute solution 
of potassium dichromate or sodium dichromate is treated with 
sulfuric acid, no reaction apparently takes place. However, if 



414 CHEMISTRY AND ITS USES 

there is present a third, substance capable of oxidation, the dichro- 
mate gives up a portion of its oxygen to this substance, and both 
the potassium and the chromium are converted into sulfates. The 
oxidation of ferrous sulfate by potassium dichromate is a good 
illustration, the reaction being represented in two steps : 

K 2 Cr 2 7 + 4 H 2 S0 4 ► K 2 S0 4 + Cr 2 (S0 4 ) 3 + 4 H 2 + 3 

6 FeS0 4 + 3 H 2 S0 4 + 30 > 3 Fe 2 (S0 4 ) 3 + 3 H 2 

This reaction is often employed in the analysis of iron ores. 

EXERCISES 

1. For what purposes are steel alloys containing manganese and 
chromium used? 

2. What four uses of manganese dioxide (pyrolusite) have been 
mentioned ? 

3. Manganese dioxide in powdered form resembles charcoal powder. 
How could you easily distinguish between the two ? 

4. Instances are on record where charcoal was mistaken for man- 
ganese dioxide and used with the intention of preparing oxygen, with 
the result that a very serious explosion took place. Explain. 

5. (a) Write the equations for the reactions which take place when 
manganese (bivalent) dissolves in hydrochloric acid ; in dilute sulfuric 
acid, (b) What are the names of the resulting salts? 

6. What is the Goldschmidt process ? 

7. When manganese dioxide is used in preparing chlorine what 
compound of manganese is formed ? 

8. How could you prepare manganous sulfide (insoluble) from 
manganous chloride (soluble) ? 

9. (a) What important sulfates have we studied that are colorless? 
(b) What ones are colored? 

10. When a solution of potassium permanganate is used as an oxi- 
dizing agent, the deep-purple color of the solution changes to a pale 
pink. Explain (see equation, p. 412). 

11. (a) What substances studied are used as disinfectants ? (b) What 
are used as antiseptics? 

12. What metal studied, in addition to manganese and chromium, 
forms acids? 



MAKGAKESE AND CHROMIUM 415 

13. (a) Write the equation for the reaction which takes place when 
chromic hydroxide is neutralized by sulfuric acid. (6) What is the 
name of the resulting salt? 

14. The common cleaning fluid used in the laboratory for cleaning 
test tubes and beakers consists of a mixture of sodium dichromate and 
sulfuric acid. Wherein does its efficiency lie? 

15. What weight of potassium permanganate could be prepared from 
100 kg. of pyrolusite, assuming that all the manganese of the ore is 
utilized? (Suggestion. 1 gram-molecular weight of the ore will form 
1 gram-molecular weight of the permanganate.) 

16. 200 lb. of ferrochromium containing 40 per cent of chromium was 
added to a ton of steel. What per cent of chromium did the resulting 
alloy contain? 



CHAPTER XLV 
PLATINUM AND GOLD 



Name 


Symbol 


Atomic 
Weight 


Density 


Melting 
Point 


Platinum 

Gold 


Pt 

Au 


195.2 
197.2 


21.50 
19.32 


1755° 
1063° 



General. Gold has been known from the earliest times and 
has always been highly prized. The Romans called it aurum, 
from which word its symbol is derived. We commonly think 
of gold as the most valuable of the precious metals, yet plati- 
num is worth about five times as much. The value of gold 
and platinum depends partly upon their relative scarcity, 
partly upon their properties, and partly because the money 
value of gold is fixed by law. Neither of them is acted upon 
by air or water. They are inactive toward most reagents and 
will not dissolve in the common acids, although readily soluble 
in aqua regia. Both are very heavy, and the melting point of 
platinum is 'high. 

Gold 

Properties. Gold is a very heavy bright-yellow metal, ex- 
ceedingly malleable and ductile, and a good conductor of elec- 
tricity. Its melting point (1063°) is much below that of 
platinum. It is quite soft and is usually alloyed with copper 
or silver to give it the hardness required for most practical 
uses. The degree of fineness is expressed in terms of carats, 
pure gold being 24 carats ; the gold used for jewelry is usu- 
ally 18 carats, 18 parts being gold and 6 parts copper or silver. 

416 



PLATINUM AXD GOLD 



417 



Gold coinage is 90 per cent gold and 10 per cent copper. Gold 
is the basis of international credit, and its price is fixed at 
$20.67 per troy ounce. 

Occurrence. South Africa produces the most gold. The 
United States produces about one fifth of the world's sup- 
ply, Alaska, California, Colorado, and Nevada leading in its 




Fig. 251. Breaking down gold-bearing deposits by streams of water 

production. Traces of it are found in almost all soils and even 
in sea water, but the cost of its production from such sources 
far exceeds the value of the metal extracted. 



Extraction. The extraction of gold is accomplished in a number 
of ways, according to the character of the deposit. In placer 
mining the gold-bearing sand is washed by a current of water 
which is so regulated that particles of light weight are swept 
away, while the heavier gold is obtained as a sediment. In hy- 
draulic mining the earth and sand are swept into sluices by pow- 
erful streams of water operated by pumps. In quartz mining the 
quartz is stamped to powder and is then washed over copper 



418 CHEMISTRY AND ITS USES 

plates, the surfaces of which have been amalgamated. The par- 
ticles of gold stick to the mercury or dissolve in it, the gold being 
recovered by distillation. In other cases, especially when the 
gold is in very fine powder or in chemical combination, chemical 
reactions are employed. In the cyanide process the gold-bearing 
material is treated with a dilute solution of sodium cyanide, with 
free access of air. The gold dissolves to form a complex cyanide, 
from which it can be precipitated by metallic zinc or by electrolysis. 
In the chlorination process the ore is treated .with chlorine, 
which converts the gold into the soluble trichloride AuCl 3 . It is 
recovered from this solution by suitable precipitants. 

Chemical conduct. Gold is not attacked by any one of the 
common acids ; aqua regia easily dissolves it, forming chlor- 
auric acid (HAuCl 4 ). Fused alkalies also attack it. Most oxi- 
dizing agents are without action upon it, and in general it is not 
an active element. 

Platinum 

Properties. Platinum is a grayish-white metal of high luster 
and is very malleable and ductile. It has a relatively high 
melting point (1755°) and is a good conductor of electricity. 
In finely divided form it has the ability to adsorb gases, espe- 
cially oxygen and hydrogen. These adsorbed gases are in a 
very active condition, resembling the nascent state, and can 
combine with each other at ordinary temperatures. A jet 
of hydrogen or coal gas directed upon spongy platinum 
quickly ignites. 

Occurrence and production. In normal times Russia furnishes 
by far the largest quantity of platinum. Thus in 1914 (pre- 
vious to the World War) that country produced 241,200 troy 
ounces out of a total world's production of 263,453 ounces. 
During the war the supply was cut off, while the demand for 
the metal for use as a catalyzer in the production of certain 
chemicals greatly increased. As a result, the price rose from 
about $45 to over $100 per troy ounce. 



PLATIXOI AND GOLD 419 

Chemical conduct. Platinum is a very inactive element chemi- 
cally and is not attacked by any of the common acids. Aqua 
regia slowly dissolves it, and it is also attacked by fused alkalies. 
It combines at higher temperatures with carbon and phospho- 
rus and forms alloys with many metals. It is readily attacked 
by chlorine but not by oxidizing agents. 

Platinum as a catalytic agent. Platinum is remarkable for its 
property of acting as a catalytic agent in a large number of chem- 
ical reactions, and mention has been made of this use of the metal 
in connection with the manufacture of sulfuric acid. AVhen desired 







% 




jH^ iOk: ft -:- 




fM 



Fig. 252. Some laboratory utensils made of platinum 

for this purpose some porous or fibrous substance, such as asbestos, 
is soaked in a solution of chloroplatinic acid and then ignited. 
The platinum compound is decomposed and the platinum deposited 
in very finely divided form. Asbestos prepared in this way is 
called platinized asbestos. The catalytic action seems to be in part 
connected with the property of adsorbing gases and rendering 
them nascent. Some other metals possess this same power, notably 
palladium, which is remarkable for its ability to adsorb hydrogen. 

Applications. Platinum is very valuable as a material for the 
manufacture of chemical utensils which are required to stand 
a high temperature or the action of strong reagents. Plati- 
num crucibles, dishes, forceps, electrodes, and similar articles 
(Fig. 252) are indispensable hi the chemical laboratory. 
A considerable percentage of the supply of the metal is used 



420 CHEMISTRY AND ITS USES 

in dentistry. In the industries platinum is used as a catalytic 
material in a number of reactions. A large fraction of the 
annual production is used for jewelry. Its use in jewelry is 
due primarily to its high cost. It is unfortunate that it should 
be so used, because of the limited supply and its importance 
in necessary industries. 

Chloroplatinic acid (H 2 PtCl 6 ). When platinum is dissolved 
in aqua regia and the solution is evaporated to dryness, orange- 
colored crystals of chloroplatinic acid (H 2 PtCl 6 ) are obtained. 
The potassium and ammonium salts of this acid are nearly 
insoluble in water and alcohol. 

EXERCISES 

1. In what respects do gold and platinum differ from the other 
metals studied? 

2. Compare the densities of platinum and silver ; of gold and copper. 

3. What is fool's gold? 

4. Platinum and silver resemble each other in appearance. How 
could you easily distinguish between them ? 

5. What are the relative costs of gold and platinum? 

6. What is the position, in the displacement series, of the metals 
that occur free in nature ? 

7. (a) What action would take place if a strip of metallic silver 
were placed in solution of gold and platinum compounds ? (b) In what 
process is advantage taken of this action? 

8. Mention three substances used for imparting a ruby color to glass. 

9. (a) What are the properties which make platinum so valuable 
for use in the manufacture of chemical utensils ? (b) Why not use gold 
in its place ? 

10. What action would take place if nitric acid were added (a) to 
gold? (&) to a gold coin? 

11. A five-dollar gold piece weighs approximately 8.334 g. What 
weight of gold does it contain? 

12. What weight of chlorauric acid can be prepared from the gold 
present in a five-dollar gold piece (see preceding problem )? 



CHAPTER XLVI 
THE STORY OF RADIUM 

Introduction. In the-spring of 1921 Madame Curie (Fig. 253), 
professor of physics at the University of Paris, paid a visit to 
the United States that attracted much attention in the daily 
press. Many honors were bestowed upon her by universities 
and scientific societies. By popular subscription the women 
of our country raised the sum of $120,000, with which was 
purchased one gram of radium in the form of radium bromide, 
— the costliest of all materials, — and on their behalf this was 
presented to Madame Curie by the President of the United 
States in the White House at Washington. Never before has 
the visit of a professor from a European country attracted 
so much attention. Why was this ? 

Discovery of radium. To answer the question given above 
we must go back some twenty-six years. In 1895 the German 
physicist Rontgen discovered a wonderful form of radiation 
capable of passing through glass tubes and acting upon a 
photographic plate much as light acts. These were the X rays, 
by means of which the surgeon takes pictures of broken bones 
in the body. 

It occurred to the French physicist Becquerel that similar 
radiations might be given off by those substances that are 
phosphorescent in the dark. Among others he tried some 
compounds of the rare metal uranium, wrapping them in 
black paper to exclude light and placing them on a photo- 
graphic plate. Immediately under these salts the plate was 
affected as by light, so there ivas a radiation of some kind. 

421 



422 CHEMISTEY AND ITS USES 

Two facts of great importance were soon discovered: (1) these 
rays could not pass through metal objects, so that radiographs 
could be made with them just as with X rays (Fig. 256) ; 
(2) when brought near these uranium compounds a charged 
electroscope was rapidly discharged (Fig. 257), showing that 
the air all around them was an electrical conductor, which is 
not true of ordinary air. The property that caused these 
effects was called radioactivity. 

Fig. 257 represents a simple form of aluminium-leaf electro- 
scope, the leaves spreading apart (indicated at B) when an electric 
charge is communicated to the knob A. When a substance con- 
taining uranium (Fig. 258, C) is brought near the knob the charge 
is rapidly lost, and the leaves collapse, as shown at B. 

Work of the Curies. At the suggestion of Becquerel his 
associate Pierre Curie and his Polish bride, Madame Curie, 
took up the problem of radioactivity. They found that pitch- 
blende, U 8 O g (the common ore of uranium), was four times as 
active as pure uranium oxide, and they argued that there 
must be present in pitchblende a radioactive substance that is 
not entirely removed in the process of purification. By long- 
continued work on tons of ore they finally isolated a minute 
quantity of the chloride of a new element four million times 
as radioactive as uranium, and which they named radium. 

In 1906 Monsieur Curie was killed in a street accident, and 
Madame Curie continued the work, in which many others now 
engaged. She found the atomic weight of radium to be 226, 
and this weight, as well as all the properties of the compounds 
of radium, placed the element in the vacant place below barium 
in the calcium family. A minute quantity of the metal itself 
was isolated by Madame Curie in 1910. 

Quantity of radium available. In all other instances when 
a new and very interesting element has been found, energetic 
search has soon provided a supply. This was not so with 
radium. It was soon found out that radium occurs in all 




Fig. 253. Madame Curie in her laboratory in Paris 




Fig. 254. Picture taken in the laboratory of the 
Standard Chemical Company, showing the crystal- 
lization of the radium bromide which the women 
of America presented to Madame Curie" on the 
occasion of her visit to America in 1921 



> PRESENTED BY TH E FEES! D £1 IT 3 F TH E Q 
UNITED STATES fJF AMERICA ON BEHALF 
OFTHE WHMI'.I! AMERICA 
Tfl 

MADAME MARIE SKLO IIOWSKA CURIE 

SERVICE TO SCIENCE AND T U HUMXNlTY 
IN f!ir (M CfiVI K ,' . i r. \j> 

THE WHITE HOUSE. VVASHfNGTTJN Of, 

> ami' j !.i N i F IN ! I..NI ^ 



". 



Fig. 255. Inscription on the box containing the radium 

bromide presented to Madame Curie by the women 

of America 



THE STORY OF RADIUM 



423 



uranium ores and in no others. A still more surprising fact is 
that in all these ores the ratio between the uranium and the radium 

is constant, being about 
1 part of radium to 
3,200,000 parts of ura- 
nium. There is no hope 
of finding any other 
of radium, and 
the in- 
very rare. 



af^V 



ores 



Fig. 256. A radiograph of some metal objects 




Fig. 257. A charged 
electroscope 



uranium ores 

selves are 

The older source was 

pitchblende, occurring chiefly in Austria. A more abundant 

source is the rare mineral carnotite, found 

in the deserts of Colorado and Utah. 

Up to the present the world's produc- 
tion of radium has been about 160 g. (say 
6 oz.), of which the United States has 
produced about 120 g. In 1920 the pro- 
duction was 35 g. It takes about 500 tons 
of ore and an equal weight of chemicals 
to produce one gram. The present value is 
about f 120,000 for every gram of radium content in a salt. 
Disintegration of radium. As one looks 
casually at a sample of a radium salt, 
such as the chloride, there is little to 
distinguish it from any other white salt, 
such as its first cousin, barium chloride. 
But if we could watch many millions of 
radium atoms we should see a wonderful 
sight. Every now and then we should 
see one of these atoms explode violently, 
much as now and then a grain of pop corn explodes as we 
warm the corn, and it then ceases to be a radium atom. 
The rate at which these explosions take place makes the 




Fig. 258. A discharged 
electroscope 



424 CHEMISTRY AND ITS USES 

average age of a radium atom about twenty-five hundred 
years. As a result of these explosions all compounds of 
radium (containing uncounted millions of atoms) are con- 
stantly giving off three kinds of so-called rays. These ar,e 
designated as alpha rays, beta rays, and yamma rays. 

A. gram of radium shoots off every second, at a velocity of 1200 
miles per second, 145,000 billion particles that are the alpha rays. 
They consist of positively charged helium atoms. At the same 
time 71,000 billion particles, constituting the beta rays, are shot off, 
and these have a velocity of about ten times that of the alpha 
rays. They are negatively charged, and in mass are about T ^-g 
as heavy as hydrogen atoms. They are now called electrons and are 
really atoms of electricity. The gamma rays are not made up of 
matter, but are vibrations in the ether like light. They correspond 
to the flash of light that often accompanies an ordinary explosion 
and move with the same velocity as light, but they differ from 
light in having very short wave-lengths like X rays. The rate of 
these explosions is kept up very steadily all the time and is not 
changed by any means as yet tried, — such as very high or very 
low temperatures, — and all compounds of radium act just alike in 
the rate of these explosions. 

Where radium comes from. Radium is decomposing at a 
rate that puts its average age at twenty-five hundred years, 
and yet the rocks in which it is found are very much older 
than this. Why has any radium survived ? Plainly it must 
be in constant process of formation from some other element, 
and it has been proved that this parent element is uranium. 
Uranium is also going to pieces (it is radioactive), but at a 
much slower rate than radium. So all uranium ores must con- 
tain radium, and none can contain very much. The quantity 
of radium present is merely the result of equilibrium between 
the rate of formation and the rate of disintegration. So we 
see why we shall never find rich radium ores. If anyone 
were to find such an ore we should have to change all our 
ideas about radium. 






THE STORY OF RADIUM 425 

Uranium series of elements. It must not be thought that an 
atom of radium or of uranium goes entirely to pieces, forming- 
nothing but alpha and beta rays, when it explodes. The loss of a 
pair of these rays by a uranium atom creates a new atom, which 
by a fresh explosion gives rise in turn to still another atom. 
This process continues until, finally, an atom is formed which does 
not explode, and all radioactivity then ceases. Th Is final element 
is lead. There are in this series thirteen radioactive elements be- 
tween uranium and lead, radium being the fifth. The immediate 
product of radium is the gaseous element niton, and its explosion 
is the most terrific of the whole series. 

Internal energy of radium. If the explosions of the radium 
atoms and of the succeeding atoms are as terrific as we have 
pictured them, and if the alpha and beta rays are shot off 
with such incredible velocities, the energy set free must be 
much greater than that of any known reaction between ele- 
ments. The radium compound actually keeps warm all the 
time. It is estimated that 1 g. of radium hourly evolves 132 cal. 
of heat and this is kept up for the average life of twenty-five 
hundred years. From this it can be computed that the total 
energy given off by a gram of radium in disintegrating into 
lead is 300,000 times the heat of combustion of a gram 
of carbon. 

Why radium is so important. The way in which the atoms 
of radium explode, the products formed in the explosion, and 
the similar conduct of all the elements of the series in which 
radium has a place make us feel certain that here is a series 
of elements, of which radium is the best known, whose atoms 
have an elaborate structure made up of helium atoms and 
electrons. These must be arranged in such a way as to hold 
an enormous store of energy — much as a compressed spring 
holds energy. The very fact that radium resembles the other 
elements in all ways, apart from radioactivity, and has its 
natural place in the periodic table makes us suspect that all 
the other elements have a similar structure, only they explode 



426 CHEMISTRY AND ITS USES 






so rarely that we never know it. In fact, we now have many 
other reasons for knowing that this must be the case. It is 
this revelation as to the nature of atoms that gives to radium 
its supreme importance as a scientific discovery. The facts 
seem to indicate that the atoms of all the elements are made 
up of unit parts and that the nature of these units is the 
same irrespective of the element. 

Practical uses of radium. Such a remarkable material as 
radium could hardly fail to find useful applications. The 
rays emitted from radium, niton, and other members of the 
series, particularly the gamma rays, produce many chemical 
and physiological effects. They disintegrate glass, water, and 
many other substances. They render certain materials lumi- 
nous in the dark, and enamel paints containing radium are 
used to illumine the hands of watches, the push buttons of 
electric lights, and the keyholes of doors. They produce 
severe burns upon the skin like those of X rays. They kill 
bacteria and other microorganisms. 

This latter property has led to the hope that exposure to 
the radiations of radium compounds might, under proper 
conditions, prove to be of service in effecting a cure for some 
diseases of the skin and for cancer. It is not possible to say 
as yet to what extent these hopes will be realized. Certain 
forms of cancer have almost certainly been cured in this way, 
and the progress of other forms has been delayed. 

Radioactive thorium. Madame Curie discovered that the rare 
element thorium exhibits properties very similar to those of 
uranium. It gives rise to a similar series of radioactive elements 
by successive decomposition, producing the same varieties of radi- 
ation as the other series. Uranium and thorium are the elements 
of greatest atomic weight, and no ordinary elements are known to 
possess similar properties. This suggests the idea that possibly 
elements of still higher atomic weight may have existed at some 
time, but that they have disintegrated into ones of smaller atomic 
weight which are not radioactive. 



THE STOKY OF RADIUM 427 

EXERCISES 

1. What is the meaning of the words alpha, beta, and gamma f 

2. What is the velocity of light (see physics) ? 

3. How is an electroscope charged? 

4. In what connection has the element thorium been mentioned? 

5. One of the radioactive elements discovered by Madame Curie was 
named polonium. Can you suggest a reason for the name? 

6. The fastest-moving projectiles from a cannon (the German big 
berthas that bombarded Paris) move a mile a second. How does the 
velocity of the alpha ray compare with this ? 

7. If abundant deposits of uranium ores should be discovered, would 
radium ever become a cheap element? 

8. Why should scientists object to the use of radium in luminous 
paints ? 

9. Would it necessarily be true that all lead was once radium? 



CHAPTER XLVII 
SOME APPLICATIONS OF RARER ELEMENTS 

Rarer elements. A large number of elements are known 
which have not been described in the foregoing pages because 
an acquaintance with them is not at all necessary for an under- 
standing of the principles of chemistry. Some of these, while 
comparatively rare, could be produced in considerable quan- 
tities if there were any commercial use for them. A good 
example is tellurium, an element in the sulfur family obtained 
as a by-product in copper refining. Others are so rare that the 
cost of production is prohibitive, even though they have very 
useful properties. 

Application in the industries. Some of these less familiar 
elements or their compounds have properties which make them 
valuable for special purposes, and mention of a few of these 
applications will be of interest. 

The rare earths constitute a group of about sixteen elements, 
all resembling aluminium in a general way. They are very 
difficult to separate from each other and always occur together 
in nature. Very large quantities of a mixture of them accumu- 
late in the extraction of thorium from monazite sand (p. 230). 
The only one whose compounds are obtained pure rather easily 
is cerium. Compounds of cerium are used as mordants, as cata- 
lytic agents, and in medicine and photography. An alloy of 
cerium with iron is used as a gas or cigar lighter, since it gives 
off a stream of sparks when scratched by hard iron. 

Thorium oxide, mixed with 1 per cent of cerium oxide, con- 
stitutes the material of which most gas mantles are made 
(p. 230). 

428 



SOME APPLICATIONS OF RARER ELEMENTS 429 



' Zirconium is found in abundance in Brazil in the form of 
the oxide Zr0 2 . This oxide has an exceedingly high melting 
point and is made into bricks, and these are used (under the 
name zirdte) for lining furnaces. 

Titanium in the silicon family is an abundant element, 
occurring chiefly as the oxide Ti0 2 , called rutile, and as a 
constituent of certain iron ores (ilmenite). Large quantities of 
nearly pure titanium or of ferrotitanium are used in making 
steel rails designed to stand very heavy wear (railway curves 
and terminals). Titanium oxide is also incorporated in electric- 
arc carbons (flaming arc), which then give a more diffused 
and efficient light than those made from pure carbon. The 
oxide is also used to impart a yellow color to porcelain and 
to artificial teeth. 

Vanadium occurs in considerable quantities in carnotite 
(p. 423) and in certain sulfides found in Peru. Ferrovanadium, 
like ferrotitanium, is used in producing spe- 
cial grades of steel, particularly when great 
toughness is desired (automobile parts). Its 
compounds are used as photographic devel- 
opers, as catalytic reagents in the dye industry 
(aniline black), as coloring materials in glass, 
and as mordants. 

Molybdenum compounds are used in col- 
oring pottery and in dyeing silk, wool, and 
leather. 

Tungsten compounds are produced hi fairly 
large quantities. It has been found possible 
to draw the metal into very fine wire (0.3 mm.), which has 
largely replaced carbon as a filament for electric lamps. Its 
melting point is very high (3400°), and the consumption of 
electrical energy for a given candle power is so low that the 
lamp is about three times as efficient as the older (carbon) 
lamp. The metal has replaced platinum for electrical contacts 




Fig. 259. The or- 
dinary type of 
tungsten lamp 



430 CHEMISTKY AND ITS USES 

in switches, telephone jacks, and automobile vibrators. Ferro- 
timgsten is used in making steel designed for lathe tools, since 
such steel can be heated to a red glow without losing temper. 

Uranium oxide is used in making the greenish-yellow glass 
so familiar in lamp shades, and as a pigment in china glazes. 

Selenium, an element in the sulfur family, is obtained as 
a by-product in refining copper. It is a nonconductor of elec- 
tricity when in the dark, but becomes a fairly good conductor 
when exposed to light. This has led to its use in automatic fire 
alarms and for regulating automatic gas buoys at sea. Added 
to glass it produces a fine red color (p. 368), such glass being 
used for automobile tail-lights and railway lanterns. It is also 
used to produce red enamels. 

Iridium gives a very hard alloy with platinum and is used for pen 
points, compass bearings, and standard weights and measures. 

Palladium is only about half as heavy as platinum, melts much 
lower, and is harder. It is used as a solder for platinum, for 
making graduated scales in scientific instruments, and as a sub- 
stitute for platinum in jewelry. In the form of a powder it is 
a remarkably active catalytic agent. 



APPENDIX 



CHEMICAL LIBRARY 

Every high school should have at least a few books dealing in 
a popular way with topics of interest to students beginning chem- 
istry. A large number of such books is available ; all are interest- 
ing, and some as fascinating as any story of romance ever written. 
The titles of a few such books are given below. Schools having 
only a limited sum of money for the purchase of books should select 
those marked with an asterisk. Prices can be obtained from any 
bookstore. A great deal of important literature bearing upon 
chemical topics may be had free of charge. This is. true of the 
bulletins published by certain departments of the Federal govern- 
ment, especially the Department of Agriculture and the Bureau 
of Standards, Washington, D.C. A list of available bulletins 
may be obtained by addressing these departments. 

Much that is interesting in regard to the subject will be found in 
the files of various periodicals. Every laboratory should subscribe 
for School Science and Mathematics, published nine times a year 
by Smith and Turton, Mount Morris, Illinois. The Independent, 
published by the Independent Corporation, New York City, con- 
tains occasional articles on chemistry by Edwin Slosson, who is 
the best of modern writers on popular chemistry. The Scientific 
American also contains numerous articles on chemical topics. 

List of Supplementary Books 

Allyn. Elementary Applied Chemistry. Ginn and Company. 
Bailey. A Textbook of Sanitary and Applied Chemistry. The 

Macmillan Company. 
Bancroft. Applied Colloid Chemistry. McGraw-Hill Book Company. 
Baskerville. Municipal Chemistry. McGraw-Hill Book Company. 
Benson. Industrial Chemistry. The Macmillan Company. 

431 



432 CHEMISTRY AND ITS USES 

*Bird. Modern Science Reader. The Macmillan Company. 
Bloxam. Inorganic and Organic Chemistry. P. Blakiston's Son & Co. 
Century Science Series. Separate biographies of Dalton, Davy, Faraday, 

Liebig, and Pasteur. 
Cushman. Chemistry and Civilization. Richard G. Badger. 
Davy. The Elementary Nature of Chlorine. The University of Chicago 

Press. 
*Duncan. The Chemistry of Commerce. Harper & Brothers. 
*Duncan. The New Knowledge. Harper & Brothers. 
*Duncan. Some Chemical Problems of Today. Harper & Brothers. 
*Faraday. Chemical History of a Candle. Harper & Brothers. 
Faraday. The Liquefaction of Gases. The University of Chicago 

Press. 
*Findlay. Chemistry in the Service of Man. Longmans, Green & Co. 
Freak. Breakfast Foods, Bulletin 162, Dairy and Food Division, Depart- 
ment of Agriculture, Harrisburg, Pennsylvania. 
Grant. Chemistry of Bread-Making. Longmans, Green & Co. 
*Green. Coal and Coal Mines. Houghton Mifflin Company. ' 
Hale. American Chemistry. D. Van Nostrand Company. 
Halligan. Fundamentals of Agriculture. D. C. Heath & Co. 
Harrow. Eminent Chemists of our Time. D. Van Nostrand Company. 
*Hart. Leavening Agents. Chemical Publishing Co. 
*Hawk. What we Eat. Harper & Brothers. 
*Hendricks. Everyman's Chemistry. Harper & Brothers. 
*Jenks. Chemistry for Young People. Frederick A. Stokes Company. 
Jordon. Principles of Human Nutrition. The Macmillan Company. 
Lassar-Cohn (translated by Muir). Chemistry in Daily Life. 

H. Grevel Co., London. 
McPherson and Henderson. A Course in General Chemistry. Ginn 

and Company. 
*Martin. The Story of a Piece of Coal. George Newnes, Ltd., London. 
*Martin. Triumphs and Wonders of Modern Chemistry. D. Van 

Nostrand Company. 
*Meade. Story of Gold. D. Appleton and Company. 
Muir. Heroes of Science — Chemists. E. & J. B. Young & Co. 
*Muir. The Story of Alchemy. D. Appleton and Company. 
Olsen. Pure Foods. Ginn and Company. 

Philip. Romance of Modern Chemistry. J. B. Lippincott Company. 
*Pilcher and Jones. What Industry owes to Chemical Science. 

D. Van Nostrand Company. 



APPENDIX • 433 

*Priestley. The Discovery of Oxygen. The University of Chicago 
Press. 

Ramsay. Gases of the Atmosphere. The Macmillan Company. 

Rose. Feeding the Family. The Macmillan Company. 

*Scheele. The Discovery of Oxygen. The University of Chicago 
Press. 

Scheele. The Early History of Chlorine. The University of Chicago 
Press. 

Scientific American Cyclopedia of Formulae. Munn & Company, New 
York City. 

Sherman. Chemistry of Food and Nutrition. The Macmillan Company. 

*Slosson. Creative Chemistry. The Century Co. 

Smith. Chemistry in America. D. Appleton and Company. 

Smith. Life of Robert Hare. J. B. Lippincott Company. 

*Smith. Story of Iron. D. Appleton and Company. 

*Sxell. Elementary Household Chemistry. The Macmillan Company. 

*Snyder. Chemistry of Plant and Animal Life (popular). 

Stewart. Chemistry and its Borderland. Longmans, Green & Co. 

Surface. Story of Sugar. D. Appleton and Company. 

*Tower. Story of Oil. D. Appleton and Company. 

*Vexable. A Short History of Chemistry. D. C. Heath & Co. 

Vulte and Venderbilt. Food Industries. Chemical Pub. Co. 

Wardell and White. A Study of Foods. Ginn and Company. 

Wiley. Foods and their Adulteration. P. Blakiston's Son & Co. 

Woolman and McGowax. Textiles. The Macmillan Company. 

*Wood. Story of a Loaf of Bread. G. P. Putnam's Sons. 

United States Department of Agriculture : (1) Composition of Foods, 
Bulletin 28, Office of Experiment Stations ; (2) Nutritive Value of 
Foods, Farmers' Bulletin 142; (3) Some Forms of Food Adulteration 
and Simple Methods for their Detection, Bulletin 100, Bureau of 
Chemistry ; (4) Industrial Alcohol, Fanners' Bulletins 268, 269 ; 
(5) Household Tests for the Detection of Oleomargarine and Reno- 
vated Butter, Farmers' Bulletin 363; (7) Canned Fruits, Preserves, 
and Jellies, Farmers' Bulletin 203. (Send to the Department of Agri- 
culture, Washington, D. C, for list of available bulletins and select 
such as may be of interest.) (8) Removal of Stains from Clothing and 
other Textiles, Farmers' Bulletin 861. 

United States Bureau of Standards : Measurements in the Home, 
Circular 55; Common Materials used in the Home, Circular 70; 
Problem of Safety, Circular 75 ; Inks, Circular 95. 



434 



CHEMISTRY AND ITS USES 



THERMOMETERS 



100 

90 

SO 



A thermometer is the well-known instrument used for measuring 
temperatures. It consists of a glass bulb joined to a thick-walled 
glass tube which has a very small but uniform bore (Fig. 260). 
There are two kinds of thermometers in common use : the 
Fahrenheit (F.), ordinarily used in our homes, and the centi- 
grade (C), used for all scientific purposes. These differ only 
in the manner in which they are graduated. In their construction 
the bulb and tube are filled with mercury, and the bulb is then 
surrounded by melting ice. The mercury con- 
^ ^ tracts as it reaches the temperature of the 

melting ice and remains stationary when this 
temperature is reached. The height of the 
mercury in the tube is then marked on the 
tube. The thermometer is next immersed in 
boiling water or steam (under a pressure of 
1 atmosphere). The height to which the mer- 
cury rises is again marked on the tube. In 
the Fahrenheit thermometer the melting point 
of ice and the boiling point of water are marked 
32° and 212° respectively, and the space 
between these two points on the tube is 
divided into 180 equal parts (degrees) ; in 
the centigrade thermometer the two points 
noted above are marked 0° and 100° respec- 
tively, and the intervening space on the ther- 
mometer tube divided into 100 equal parts 
(Fig. 260). In other words, in the Fahren- 
heit system ice melts at 32° and water boils 
at 212°, while in the centigrade ice melts at 0° 
and water boils at 100°. In the construction of thermometers 
the divisions on the tube may be extended below the melting 
point of ice and above the boiling point of water and, of course, 
must be of the same length. 

It is easily possible to convert readings on the Fahrenheit scale 
into the corresponding readings on the centigrade scale and vice 
versa. It is only necessary to keep in mind that 180° on the 
Fahrenheit scale (the difference in degrees between the melting 



I 



212" 

194" 
176° 
158° 
140° 
122" 
104° 
86° 
68° 
50° 



Fig. 260, The cen- 
tigrade and Fahren- 
heit scales 



APPENDIX 435 

point of ice and the boiling point of water) equals 100° on the 
centigrade (or 1° F. equals |° C.) and that 32° F. is the same as 
0° C. The above relations are expressed in the following equation : 

C 5 



F-32 9 

Suppose we wish to convert 75° F. into centigrade reading : sub- 
stituting 75 for F in the above equation, we have 

C 5 



75-32 9 

Solving for C, we have 9 C = 5 (75 - 32) 

9 C = 215 
C = 23.8° F. 

Similarly, if we wish to convert centigrade readings into Fahren- 
heit we substitute the centigrade reading for C in the above 
equations and solve for F. 



436 CHEMISTRY AND ITS USES 

DENSITIES AND MELTING POINTS OF SOME COMMON ELEMEN rs 



Name Density 

Aluminium . . . 2.65 
Antimony .... 6.70 

Arsenic 5.73 

Bismuth .... 9.80 
Calcium .... 1.55 
Carbon (diamond) 3.52 
Carbon (graphite) . 2.30 
Chromium . . . 7.30 

Cobalt 8.60 

Copper 8.93 

Gold 19.32 

Iron 7.86 

Lead 11.37 

Magnesium . . . 1.74 






Melting 






Melting 


Point 


Name 


Density 


Point 


658.7° 


Manganese . . 


. 8.01 


1230.0° 


630.0° 


Mercury . . . 


. 13.56 


-38.87° 





Nickel .... 


. 8.90 


1452.0° 


271.0° 


Phosphorus . . 


. 1.83 


44.0° 


810.0° 


Platinum . . . 


. 21.50 


1755.0° 





Potassium . . . 


. 0.862 


62.3° 


3600.0° 


Radium .... 


. 


700.0° 


1615.0° 


Silicon .... 


. 2.35 


1420.0° 


1480.0° 


Silver .... 


. 10.50 


960.5° 


1083.0° 


Sodium .... 


. 0.97 


97.5° 


1063.0° 


Sulfur (rhombic) 


. 2.06 


112.8° 


1530.0° 


Tin 


. 7.30 


231.9° 


327.4° 


Zinc 


. 7.10 


419.4° 


651.0° 









TABLE OF SOLUBILITY OF VARIOUS SOLIDS 







Weight dissolved by 100 cc 


of Water at 


Substance 












0° 


20° 


100° 


Calcium chloride . . 


CaCl 2 


59.5 g. 


74.5 g. 


159.0 g. 


Sodium chloride . . . 


NaCl 


35.70 g. 


36.0 g. 


39.80 g. 


Potassium nitrate . . 


KN0 3 


13.30 g. 


31.6 g. 


246.0 g. 


Copper sulfate . . . 


CuS0 4 


14.30 g. 


20.7 g. 


75.4 g. 


Calcium sulfate . . . 


CaS0 4 


0.759g. 


0.203 g. 


0.162 g. 


Calcium hydroxide . . 


Ca(OH) 2 


0.185 g. 


0.165 g. 


0.077 g. 



TENSION OF AQUEOUS VAPOR EXPRESSED IN MILLIMETERS 
OF MERCURY 



Temperature Pressure 

15° 12.78 

16° 13.62 

17° 14.52 

18° 15.46 

19° 16.56 



Temperature Pressure 

21° 18.62 

22° 19.79 

23° 21.02 

24° 22.32 

25° 23.69 



20° 



17.51 100° 760.00 



INDEX 



Acetanilide, 236 

Acetates, 259 

Acetic acid, 106, 259 ; glacial, 259 

Acetylene, 215 

Acid anhydrides, 173 

Acid-forming elements, 301 

Acids, 145 ; binary, 147 ; character- 
istic properties of, 145 ; dibasic, 
194 ; fatty, 259 ; hydro-, 147 ; and 
ionization, 159 ; monobasic, 194 
naming of, 147; organic, 259 
strength of, 160 ; ternary, 147 
tribasic, 194 

Adsorb, 108 

Affinity, chemical, 9 

Agate, 287 

Air, 122 ; analysis of, 123 ; changes in 
composition of, 124 ; composition 
of, 122 ; composition of constant, 
126 ; constituents of, essential to 
life, 123 ; impure, 127 ; injurious 
effect of impure, 127 ; liquefac- 
tion of, 53 ; liquid, 128 ; a mix- 
ture, 126 ; sulfur dioxide in, 123 ; 
water vapor in, 123 

Alchemists, the, 2 

Alcohol, absolute, 253 ; denatured, 
254 ; ethyl, 253 ; grain, 253 ; 
methyl, 252 ; wood, 106, 252 

Alcohols, 252 

Alizarin, 235 

Alkali metals, 305 

Alkalies, 153, 305 

Alkaloids, 251 

Allotropic forms, 59, 275 

Alloys, 282 

Aluminium, 353 ; hydrolysis of salts 
of, 360 

Aluminium hydroxide, 358 



Aluminium oxide, 357 

Aluminium silicates, 368 

Alums, 360 

Alundum, 358 

Amethyst, 287 ; oriental, 357 

Ammonia, 163 

Ammonia water, 166 

Ammoniacal liquor, 219 

Ammonium compounds, 314 

Ammonium hydroxide, 166 

Ammonium radical, 164, 167 

Ammonium salts, 167, 314 

Amorphous matter, 54 

Analysis, 68 

Anhydride, sulfuric, 188, 190; sul- 
furous, 188 

Anhydrides of acids,' 173 

Anhydrous compounds, 67 

Anhydrous salts, 291 

Aniline, 234 

Aniline dyes, 235 

Anode, 134 

Anthracene, 235 

Antimony, 281 

Apatite, 276 

Aqua ammonia, 166 

Aqua regia, 171 

Argol, 261 

Argon, 80 

Arrhenius, Svante (portrait), 157 

Arsenic, 280 ; compounds of, 280 

Asbestos, 345 

Asphalt, 213 

Aspirin, 234, 236 

Atmosphere, 122 

Atomic weights, accurate determina- 
tion of, 88 ; relative, 83 ; selection 
of, from combining weights, 87 ; 
standard for, 85 



437 



438 



CHEMISTRY AND ITS USES 



Atoms, 82 

Automobile casings, facing pp. 349, 

350 
Avogadro's principle, 47 
Azote, 76 

Babbitt metal, 283 

Bakelite, 234 

Baking, aerating agents used in, 361 

Baking powders, 361 ; reactions of , 362 

Barite, 336 

Barium, 335 ; compounds of, 336 

Base-forming elements, 301 

Bases, 153 ; and ionization, 159 ; 

strength of, 160 
Bauxite, 354 
Beehive coke oven, 220 
Benzene, 234 
Benzine, 212 
Benzoic acid, 234 
Bessemer process, 378 
Bismuth, 281 ; compounds of, 282 
Bisque, 369 
Bivalent atoms, 118 
Blast furnace, 375 
Blast lamp, 37 
Bleach, 332 
Bleaching, 141 ; by chlorine, 140 ; 

by hydrogen peroxide, 73 ; by 

sulfurous acid, 189 
Bleaching powder, 138, 332 
Blue-printing, 384 
Bluestone, 393 
Boiler scale, 346 
Boiling point, 52 
Boneblack, 107 
Borax, 291 

Bordeaux mixture, 393 
Boric acid, 290 
Boron, 290 

Boyle, Robert (portrait), 41 
Brass, 392 
Brazing, 291, 402 
Bread, 256 
Brimstone, 183 



Bromides, 207 

Bromine, 206 

Bromine in the World War, 207 

Bronze, 392 ; aluminium, 355, 392 

Bunsen burners, 228 

Burning, 20 

Butter fat, 264 

Butyrin, 264 

By-product coke oven, 220 

By-products, 19, 308 

Cadmium, 350 

Caffeine, 250, 255 

Calcite, 328 

Calcium, 327 ; other salts of, 335 

Calcium acid carbonate, 328 

Calcium carbide, 334 

Calcium carbonate, 328 ; effect of 
heat on, 177 

Calcium cyanamide, 334 

Calcium family, 327 

Calcium hydroxide, 331 

Calcium oxide, 330 

Calcium phosphate, 335 

Calcium sulfate, 333 

Calomel, 395 

Calorie, 57 

Calorific value of fuels, 225 

Calorimeter, 57 ; bomb, 225 ; respira- 
tion, 272 

Candles, 321 

Caramel, 241 

Carats, 416 

Carbides, 109 

Carbohydrates, 238 

Carbolic acid, 234 

Carbon, 102 ; amorphous, 105 ; chem- 
ical conduct of, 109; crystalline, 
102 ; determination of, in com- 
pounds, 109 ; properties of, 107 ; 
retort, 220 ; uses of, 109 

Carbon dioxide, 110 ; determination 
of, in air, 124 

Carbon disulfide, 195 

Carbon monoxide, 114 



f 



INDEX 



439 



' Carbon tetrachloride, 214 

Carbona, 215 

Carbonates, solubility of, 304 

Carbonblack, 107 

Carbonic acid, 113 

Carbonic acid gas, 110 

Carborundum, 285 

Carnallite, 311, 344 

Carnotite, 423 

Casein, 241, 242 ; coagulation of , 295 

Cassiterite, 401 

Catalysis, 17 

Catalytic action of manganese diox- 
ide, 17 ; of platinum, 169, 190, 419 

Catalytic agent, 17 

Cathode, 134 

Caustic potash, 313 

Caves, formation of, 329 

Cell, Castner, 151 ; electrolytic, 138 

Celluloid, 246 

Cellulose, 246 

Cellulose acetate, 246 

Cement, 370 ; Portland, 370 

Cementite, 373 

Cerium, 428 

Chalk, French, 346; precipitated, 328 

Chamber process for sulfuric acid, 192 

Charcoal, 106 ; activated, 108 ; ani- 
mal, 107 

Cheese, 242 

Chemical action, 7 

Chemical affinity, 9 

Chemical changes, 7 

Chemical library, 431 

Chemist, work of the, 4 

Chemistry, future of, 5 ; general 
field of, 1 ; relation of, to other 
sciences, 1 

Chile saltpeter, 309 

Chlorauric acid, 418 

Chlorides, 136, 139 ; properties of, 
148 ; solubility of, 303 

Chlorine, 136 ; action of, as disinfect- 
ant, 141 ; action of, on elements, 
138 ; action of, on hydrogen, 139 ; 



action of, on water, 139 ; bleach- 
ing action of, 140 ; uses of, 141 ; 
in the World War, 142 

Chlorine family, 203 

Chlorine plant at Edgewood, facing 
p. 143 

Chloroform, 214 

Chloroplatinic acid, 420 

Chromates, 413 

Chrome iron ore, 412 

Chrome yellow, 407, 413 

Chromic acid, 413 

Chromite, 412 

Chromium, 412 ; compounds of, 413 

Cider, 260 

Cinnabar, 395 

Citric acid, 262 

Clay, 368 ; coagulation of, 296 

Clay products, 368 

Coal, 105 

Coal gas, 219 

Coal mine, 105 

Coal tar, 232 

Coal-tar compounds, 232 ; in foods, 
236 ; in medicines, 236 

Coal-tar dyes, 235 

Cobalt, 384 ; compounds of, 386 

Cocaine, 250 

Coin, gold, 392, 417 ; silver, 392 

Coinage, nickel, 385, 392 

Coke, 106, 220 

Coke oven, 220 

Cold-storage plant, facing p. 167 

Collodion, 246 

Colloidal dispersion, 293 

Colloidal state, 293 

Colloids, 293 ; coagulation of, 295 ; 
and crystallization, 297 ; dispersal 
of, 296 ; preparation of, 294 ; prop- 
erties of, 294 

Combining weights, 84 

Combustion, 21 ; and concentration, 
23 ; heat of, 24 ; products of, 22 ; 
speed of, 23 ; spontaneous, 25 ; 
and temperature, 23 



440 



CHEMISTRY AND ITS USES 



Common salt, mining of , facing p. 142 

Completion of reactions in solution, 
178 

Composition, atomic, 91 ; percent- 
age, 91 

Compounds, definition of, 7 ; number 
of, 11 

Concrete, 371 

Condensite, 234 

Contact process for sulfuric acid, 192 

Copper, 390 ; action of nitric acid 
on, 170 ; action of sulfuric acid on, 
193 ; blister, 391 ; metallurgy of, 
391 ; ores of, 391 ; refining of, 392 

Copper oxide, reduction of, by hydro- 
gen, 36 

Copperas, 382 

Corn oil, 245 

Corn sirup, 243 

Corrosive sublimate, 395 

Corundum, 357 

Cotton fiber, 249 

Cottonseed oil, 263 

Cottrell process, 299 

Coumarin, 236 

Cracking of oils, 213 

Cream of tartar, 261 

Cresylic acid, 235 

Cryolite, 354 

Crystals, 54 

Cupric compounds, 393 

Cuprous compounds, 393 

Curie, Madame (photograph), facing 
p. 422 

Cyanamide, 334 

Cyanogen, 174 

Dal ton, John (portrait), facing p. 72 
Davy, Sir Humphry (portrait), 150 
Dehydrating agent, 194 
Deliquescent compounds, 314 
Denaturants, 255 
Density, 50 
Dewar flask, 54 
Dextrin, 243 



Dextrose, 240, 243 
Diamonds, 103 ; artificial, 103 
Diastase, 243, 256 
Dichlorethyl sulfide, 217 
Dichromic acid, 413 
Displacement series, 160 
Distillation, 63 ; destructive, 107 

fractional, 213 
Dolomite, 345 
Dryer for paints, 407 
Duraluminum, 356 
Dust explosions, 24 
Dyeing, process of, 362 
Dyes, 362 
Dynamite, 324 ; gelatin, 324 

Earths, the rare, 428 

Effervescence, 113, 133 

Eggs, preservation of, 290 

Electric furnace, 226 

Electrochemical industries, 302 

Electrochemical series, 160 

Electrodes, 134 

Electrolysis, 18, 134 ; and ionization, 
158 

Electrolyte, 156 

Electroplating, with nickel, 385 ; 
with silver, 396 

Electroscope, 423 

Electrotyping, 392 

Elements, definition of, 6 ; in human 
body, 9 ; molecular weight of, 89 ; 
names of, 10 ; number of, 9 ; oc- 
currence of, 10 

Emery, 357 

Emulsifying agents, 297 

Emulsions, 297 

Energy, 54 ; chemical, 57 ; conserva- 
tion of, 55 ; transformation of, 55 

Enzymes, 256 

Epsom salts, 345 

Equations, 94 ; molecular, 94 ; prob- 
lems based on, 98 ; steps in writ- 
ing, 96 

Equilibrium, 176 ; in solution, 178 






INDEX 



441 



Esters, 262 
Etching of glass, 205 
Ether, 253 
Ethylene, 217 
Eudiometer, 69, 71 
Explosives, 169, 322 

Fabrikoid, 246 

Families, periodic, 200 

Fats, 262 

Feldspar, 289 

Fermentation, acetic, 260 ; alcoholic, 
253 ; lactic, 241 

Ferric salts, 382 ; reduction of, 383 

Ferromanganese, 411 

Ferrotitanium, 429 

Ferrous salts, 382 ; oxidation of, 383 

Ferrovanadium, 429 

Fertilizers, 340 ; sources of, 341 

Fillers, mechanical, 66 ; sand, 65 

Fire damp, 213 

Fire extinguisher, 113 

Flame reactions, 315 

Flames, 226 

Fluorine, 203 

Flux, 374 

Foams, 299 

Fogs, 299 

Foods, 266 ; fuel value of, 270 

Fool's gold, 382 

Formaldehyde, 252 

Formalin, 253 

Formic acid, 114 

Formulas, 92 ; percentage composi- 
tion from, 93 

Freezing point, 51 

Fuel gases, composition of, 223 

Fuels, 219 ; calorific value of, 225 

Fuller's earth, 368 

Galenite, 405 
Galvanized iron, 348 
Garbage, utilization of, 322 
Gas, coal, 219 ; enriched, 222 ; natu- 
ral, 223 ; producer, 222 ; water, 222 



Gases, collection of, 18 ; effect of 
pressure on, 41 ; effect of tempera- 
ture on, 42 ; liquefaction of, 52 ; 
nature of, 47 ; poison, 108 ; solu- 
bility of, 133 ; weight of a liter 
of, 89 

Gas mantles, 229 

Gas masks, 108 

Gasoline, 212 

Gay-Lussac (portrait), 44 

Gems, artificial, 358 

German silver, 392 

Glass, 366 ; color of, 367 ; varieties 
of, 367 

Glauber's salt, 307 

Glucose, 243 

Glycerin, 262, 322 

Glyceryl radical, 262 

Gold, 416 ; in jewelry, 416 

Goldschmidt reduction process, 356 

Gram-atomic weights, 94 

Gram-molecular weights, 94 

Graphite, 104 

Grease spots, removal of, 388 

Grignard, Victor (portrait), facing 
p. 309 

Guncotton, 246 

Gunpowder, smokeless, 323 

Gypsum, 333 

Haber process, 165 

Hall, Charles M. (portrait), 353 

Hall process, 353 

Halogens, 203 

Hard water, 329 

Hare, Robert (portrait), 37 

Heat, measurement of, 56 

Heat of decomposition, 67 

Heat of formation, 67, 98 

Heat of fusion, 51 

Heat of reaction, equations for, 98 

Heat of solidification, 51 

Helium, 80 ; extraction from natural 

gas, facing p. 81 
Hematite, 374 



442 



CHEMISTRY AND ITS USES 



Humus, 340 

Hydrates, 67, 291 

Hydriodic acid, 208 

Hydrobromic acid, 207 

Hydrocarbons, 211 

Hydrochloric acid, 144 

Hydrocyanic acid, 174 

Hydrofluoric acid, 205 

Hydrogen, 30; preparation of, from 
acetylene, 215 ; preparation of, 
from acids, 32 ; preparation of, 
from water, 31 ; uses of, 37 

Hydrogen bromide, 207 

Hydrogen chloride, 142 ; solubility 
of, 144 

Hydrogen cyanide, 174 

Hydrogen fluoride, 205 

Hydrogen iodide, 208 

Hydrogen nitrate, 167 

Hydrogen peroxide, 72 

Hydrogen sulfate, 191 

Hydrogen sulfide, 185 ; in air, 123 

Hydrolysis of salts, 308 

Hydrosulfuric acid, 186 

Hydroxides, solubility of, 303 

Ice, manufacture of, 166 

Iceland spar, 328 

Incendiary bullets, 278 

Indicators, 154 

Indigo, 235 

Industries, chemical, 4 

Infusorial earth, 288 

Inks, 386 

Insecticides, arsenic, 281 

Iodine, 207 ; tincture of, 208 

Iodoform, 215 

Ionization and acids, 159 ; and bases, 

159 ; and neutralization, 159 ; and 

salts, 159 ; theory of, 156 
Ions, 157 
Iridium, 430 
Iron, 372 ; action of, on steam, 32 ; 

cast, 375 ; compounds of, 381 ; 

wrought, 377 



Iron ore, 374 

Isomeric compounds, 238 

Javelle water, 388 
Jellies, 298 

Kaolin, 289 

Kaolinite, 368 

Kelp, 208 

Kerosene, 212 

Kindling temperature, 25 

Krypton, 80 

Lachrymators, 207 

Lactic acid, 241 

Lactose, 241 

Lakes, 363, 408 

Lampblack, 107 

Laughing gas, 1 72 

Lavoisier (portrait), frontispiece 

Law, of Boyle, 41 ; of Charles, 44 ; 

of the conservation of energy, 56 ; 

of the conservation of matter, 50 ; 

of definite composition, 71 ; of 

Gay-Lussac, 44, 208 ; of Henry, 

133 ; of multiple proportion, 73 ; 

the periodic, 197, 200 ■; of volumes, 

of Gay-Lussac, 208 
Lead, 403 ; compounds of, 404 ; list 

of compounds of, 407 ; sugar of, 

407 ; white, 405 
Leather industry, facing pp. 335, 336 
Leblanc, Nicholas (portrait), facing 

p. 308 
Leblanc process, 307 
Levulose, 240 
Lime, 330 ; air-slaked, 331 ; chloride of, 

138, 332 ; hydrated,331 ; slaked,331 
Limekilns, 330 
Limestone, 328 
Limestone dolomitic, 345 
Lime-sulfur spray, 184 
Litharge, 404 
Lithopone, 349 
Litmus, 144 



INDEX 



443 



Lubricating oils, 212 
Lunar caustic, 397 
Lye, 152 

Magnalium, 356 

Magnesia, 345 

Magnesite, 345 

Magnesium, 344 ; compounds of, 345 

Magnesium family, 344 

Magnetite, 374 

Malic acid, 261 

Malt, 243 

Maltose, 243 

Manganese, 411 ; compounds of, 412 

Manganous salts, 412 

Marble, 328 

Mass, 50 

Mass action, 177 

Matches, 278 

Materials, developed by plants and 
animals, 3 ; supplied by nature, 3 

Matter, allotropic forms of, 59 ; amor- 
phous, 54 ; classification of, 6 ; 
conservation of, 50 ; crystalline, 
54 ; definition of, 6 ; states of, 50 

Meerschaum, 346 

Melting point, 51 

Mendeteeff (portrait), 198 

Mercerized cotton, 247 

Mercuric compounds, 395 

Mercurous compounds, 395 

Mercury, 394 

Metallurgy, 301 

Metals, 150, 301 ; compounds of, 301 

Methane, 213 ; halogen derivatives 
of, 214 

Methane series, 211 

Milk, pasteurization of, 242 ; sour- 
ing of, 241 

Minium, 405 

Mixtures, 11 

Moissan, Henri (portrait), 204 

Molasses, 240 

Molecular weights, 82 ; of the ele- 
ments, 89 ; from formulas, 93 ; 



and percentage composition, 93 ; 
relative, 83 ; standard for, 84 

Molecules, 47, 82 

Monazite sand, 230 

Mordants, 363 

Morley, E. W. (portrait), facing p. 73 

Morphine, 250 

Mortar, 332 

Moth balls, 235 

Muriatic acid, 144 

Mustard gas, 217 

Naphtha, 212 

Naphthalene, 235 

Nascent state, 142 • 

Native state, 10 

Natural gas, 223 

Negative, photographic, 398 

Neon, 80 

Neutralization, 154 ; balancing equa- 
tions of, 154 ; and ionization, 159 

Nickel, 384 ; compounds of, 386 

Nicotine, 250 

Niton, 425 

Nitrates, 171 ; solubility of, 303 

Nitric acid, 167 ; action of, on metals, 
170; salts of, 171 

Nitric oxide, 172 

Nitrides, 79 

Nitrifying organisms, 79 

Nitrobenzene, 234 

Nitrocellulose, 246, 323 

Nitro-explosives, 322 

Nitrogen, 76 ; acids of, 167 ; com- 
pounds of, 163 ; determination of, 
in air, 124 ; oxides of, 172 ; prep- 
aration of, 77 ; utilization of, 342 

Nitrogen dioxide, 173 

Nitrogen tetroxide, 173 

Nitroglycerin, 322, 324 

Nitrous acid, 171 

Nitrous oxide, 172 

Nonmetals, 150 

Oil, lubricating, 212 
Oil dag, 296 



444 



CHEMISTRY AND ITS USES 



Oil.'of vitriol, 191 

Oil wells, facing p. 212 

Oils, 262 ; hydrogenation of, 263 

Oleic acid, 262 

Olein, 262 

Oleomargarine, 264 

Opal, 287 

Open-hearth process, 379 

Organic acids, 259 

Oxalic acid, 261 

Oxidation, 21, 383 ; heat of, 24 

Oxides, 22 

Oxidizing agent, 37 

Oxyacetylene blowpipe, 216 

Oxygen, 14 ; chemical conduct of, 
19 ; commercial preparation of, 
19 ; determination of, in air, 123 ; 
importance of, 26 ; preparation of, 
in the laboratory, 17 ; preparation 
of, from mercuric oxide, 16 ; prep- 
aration of, from potassium chlo- 
rate, 16; preparation of, from 
water, 18 ; a standard for atomic 
weights, 85 

Oxyhydrogen blowpipe, 37 

Ozone, 58 ; in air, 123 

Paints, 407 

Palladium, 419, 430 

Palmitic acid, 259, 262 

Palmitin, 262 

Paper, 249 

Paraffin, 212 

Paris green, 281 

Pectin, 298 

Pepsin, 256 

Periodic group, 200 

Periodic law, the, 197 ; value of, 201 

Periodic table of elements, 199 

Perkin, William H. (portrait), facing 

p. 235 
Permanganic acid, 412 
Petroleum, 212 
Pewter, 402 
Phenol, 234 



Phlogiston, 20 

Phosgene, 26, 115 

Phosphate fertilizers, 341 

Phosphates, 279 

Phosphorescence, 277 

Phosphorite, 276 

Phosphorus, 275 ; compounds of , 277 ; 
slow combustion of, 59 ; uses of, 
278 

Photography, 397 

Physical changes, 8 

Picric acid, 234, 323 

Pigments, 408 

Pitch, 237 

Pitchblende, 422 

Placer mining, 417 

Plant food, 340 

Plaster, 332 

Plaster of Paris, 333 

Platinized asbestos, 419 

Platinum, 418 ; as a catalyzer, 419 

Pneumatic trough, 18 

Poison gas, 115 

Pop, 255 

Porcelain, 369 

Potash fertilizers, 342 

Potassium, 311 ; compounds of, 313 

Potassium chlorate, 312 ; decompo- 
sition of, 95 ; derivation of for- 
mula of, 92 

Potassium dichromate, 413 

Potassium f erricyanide, 384 

Potassium f errocyanide, 383 

Potassium iodide, 208 

Potassium nitrate, 312 

Potassium permanganate, 412 

Potassium salts, source of, in the 
United States, 313 

Pottery, 369 

Precipitate, 179 

Precipitation, 302 

Preservatives, 256 

Pressure, standard, 42 

Priestley, Joseph (portrait), facing 
p. 16 



INDEX 



445 



Problems, solution of, 26, 38, 115 
Producer gas, 222 
Proofs, photographic, 398 
Properties, definition of, 11 
Protein, 76 
Protein matter, 163 
Proteins, 264 
Prussian blue, 384 

Prussiate of potash, red, 384 ; yel- 
low, 383 
Prussic acid, 174 
Puddling furnace, 377 
Pure Food and Drugs Act, 274 
Pyrene, 215 
Pyrex glass, 367 
Pyridine, 255 
Pyrite, 382 

Quadrivalent atoms, 119 
Quartz, 287 
Quicklime, 330 
Quicksilver, 394 
Quinine, 250 

Radiation of radium, 421 

Radical, ammonium, 164, 167 

Radicals, 146 

Radioactivity, 422 

Radium, 421 ; disintegration of, 423 ; 
importance of, 425 ; internal en- 
ergy of, 425 ; quantity of available, 
422 ; source of, 424 ; uses of, 426 

Ramsay, Sir William (portrait), fac- 
ing p. 80 

Reactions, completion of, in solution, 
178; reversible, 176 

Reducing agent, 37 

Reduction, 36, 383 ; relation of, to 
oxidation, 37 

Remsen, Ira (portrait), facing p. 4 

Rennet, 242 

Reversible reactions, 176 

Richards, T. W. (portrait), facing 
p. 73 

Rock phosphate, 335 



Rubber, 350 
Ruby, 357 
Rutile, 429 

Saccharine, 234, 236 

Safety lamp, 214 

Sal ammoniac, 314 

Sal soda, 308 

Salt, common, 306 ; rock, 306 

Saltpeter, 312 ; Chile, 151, 309 

Salts, 145 ; acid, 194 ; action of sul- 
furic acid on, 193 ; formulas of, 
146 ; hydrolysis of, 308 ; insoluble, 
302 ; and ionization, 159 ; naming 
of, 148 ; normal, 194 ; preparation 
of, 302; soluble, 302; solubility 
of, 303 

Sand, 287 

Sandstone, 287 

Saponification, 319 

Sapphire, 357 

Scheele, Karl W. (portrait), facing 
p. 80 

Scheele's green, 281 

Selenium, 430 

Serpentine, 346 

Sewage-disposal plant, 25 

Silica, 287 

Silicate industries, 366 

Silicates, 289 ; commercial applica- 
tion of, 366 

Silicic acids, 289 

Silicides, 285 

Silicon, 285 ; acids of, 289 

Silicon dioxide, 287 

Silicon tetrafluoride, 205 

Silk, artificial, 248 ; natural, 246 

Silk fiber, 249 

Silver, 396 ; compounds of, 397 

Slag, 375 

Slaking,' 331 

Smith, Edgar F. (portrait), facing p. 4 

Smoke, 229, 299 

Smoke prevention, 229 

Smokeless powder, 246 



446 



CHEMISTRY AND ITS USES 



Soap, 318 ; cleansing action of, 321 ; 
varieties of, 320 

Soapstone, 345 

Soda, 307 ; baking, 309 ; bicarbonate 
of, 309 ; caustic, 152 ; sal, 308 ; 
washing, 308 

Soda ash, 307 

Sodium, 150 ; action of, on water, 
31 ; compounds of, 152, 305, 310 

Sodium benzoate, 234, 236, 257 

Sodium carbonate, 307 

Sodium chloride, 306 ; electrolysis 
of, 138, 158 

Sodium cyanide, 174 

Sodium dichromate, 413 

Sodium family, 305 

Sodium ferrocyanide, 383 

Sodium hydrogen carbonate, 309 

Sodium hydroxide, 152 ; electrolysis 
of, 151 

Sodium nitrate, 309 

Sodium phosphate, 278 

Sodium sulfate, 307 

Sodium tetraborate, 290 

Softening of water, 329 

Soils, 338 ; kinds of, 339 ; testing of, 
for fertilizers, 342 

Solder, 402 

Soldering, 291, 402 

Solubility, 131 ; causes affecting, 132 ; 
of salts, 303 ; effect of pressure on, 
133 ; effect of temperature on, 132 ; 
tables of, 133 

Solute, 130 

Solutions, 130 ; boiling point of, 133 ; 
chemical, 130 ; classes of, 132 ; 
electrolysis of, 134 ; freezing point 
of, 134 ; properties of, 133 ; sat- 
urated, 130 ; supersaturated, 131 

Solvay, Ernest (portrait), facing 
p. 309 

Solvay process, 308 

Solvent, 130 ; effect of solute on 
boiling point of, 133 ; effect of 
solute on freezing point of, 134 



Sphalerite, 349 

Spiegel iron, 411 

Spray, lime-sulfur, 184 

Stains, removal of, from textiles, 
387 

Stalactites, 329 

Stalagmites, 329 

Standard conditions, 45 

Stannic compounds, 402 

Stannous compounds, 402 

Starch, 244 

Stassfurt salts, 311 

Stearic acid, 262 

Stearin, 262 

Steel, 378 ; alloys of, 381 ; elec- 
trothermal metallurgy of, 380 ; 
hardening of, 380 ; nickel, 385 ; 
tempering of, 380 

Stellite, 385 

Stibnite, 281 

Strontium, 335 

Strychnine, 250 

Substances, 11 

Sucrose, 239 

Sugar, 239 ; grape, 243 ; invert, 240 ; 
milk, 241 

Sugar of lead, 259 

Sulfates, 194 ; solubility of, 303 

Sulfides, 187; solubility of, 303; 
uses of, in analysis, 187 

Sulfites, 190 

Sulfur, 181 ; flowers of, 183 ; in 
foods, 182 ; monoclinic, 183 ; plas- 
tic, 184 ; prismatic, 183 ; rhombic, 
183 ; roll, 183 

Sulfur dioxide, 188 

Sulfur trioxide, 190 

Sulfur water, 187 

Sulfuric acid, 191 ; action of, on 
metals, 193 ; action of, on salts, 
193 ; salts of, 194 

Sulfurous acid, 189 ; salts of, 190 

Sylvite, 311 

Symbols, 10 

Synthesis, 68 



INDEX 



447 



Talc, 345 

Tanning, following p. 335 

Tartaric acid, 261 

Tellurium, 428 

Temperature, absolute scale of, 43 

Tervalent atoms, 119 

Textile fibers, 249 

Theory, the atomic, 83 ; of ioniza- 
tion, 156 ; the kinetic, 47 

Thermite, 357 

Thermite welding, 356 

Thermometers, 434 

Thermos bottle, 54 

Thorium, 426, 428 

Tin, 401 ; block, 401 ; compounds of, 
402 

Tin plate, 402 

Tinstone, 401 

Titanium, 429 

T.N.T., 325 

Toluene, 234 

Topaz, 357 

Trinitrophenol, 325 

Trinitrotoluene, 234, 325 

Tungsten, 429 

Turpentine, 408, 410 

Type metal, 283 

Undercooling, 51 

Univalent atoms, 118 

Uranium, 421 

Uranium oxide, 430 

Uranium series of elements, 425 

Urea, 269 



Vaseline, 212 

Ventilation, 127, 225 

Vermilion, 394 

Vinegar, 260 ; malt, 260 ; mother of, 

260 ; sugar, 260 
Vitamines, 266 
Vitriol, blue, 393 ; green, 382 ; oil 

of, 191 ; white, 349 
Vulcanization, 351 

Water, 61 ; chemical conduct of, 66 ; 
composition of, by weight, 68 ; com- 
position of, by volume, 71 ; deriva- 
tion of formula of, 92 ; distilled, 
63 ; electrolysis of, 18 ; hard, 62 ; 
heat of formation of, 98 ; heat of 
fusion of, 57 ; heat of vaporization 
of, 57 ; mineral matter in, 62 ; or- 
ganic matter in, 62 ; purification of, 
63, 359; self -purification of, 66; 
soft, 62 ; vapor pressure of, 45 

Water of crystallization, 291 

Water dag, 296 

Water gas, 222 

Water glass, 289 

Water and health, 62 

Water vapor in air, determination 
of, 124 

Whey, 241 

Wood, preservation of, 349 

Wood's metal, 283 

Wool fiber, 249 

Xenon, 80 



Valence, 118 ; applications of, 119 ; 
and the periodic law, 201 ; of radi- 
cals, 146 ; standard of, 118 ; table 
of, 120 

Vanadium, 429 

Vanillin, 236 

Vapor pressure of water, 45 

Vaporization, 52 

Varnish, 409 



Yeast, 253 

Zeppelin dirigible, facing p. 17 
Zinc, 347 ; compounds of, 349 ; 

metallurgy of , 347 ; ores of, 347 
Zinc white, 349 
Zircite, 429 
Zirconium, 429 
Zymase, 256 



LIST OF THE ELEMENTS, THEIR SYMBOLS 
AND TH^IR ATOMIC WEIGHTS 

The more important elements are printed in heavier type 
= 16 



Aluminium 
Antimony 

Argon . 

Arsenic 

Barium . 

Bismuth 

Boron . 

Bromine 

Cadmium 

Caesium 

Calcium 

Carbon . 

Cerium . 

Chlorine 

Chromium 

Cobalt . 

Columbium 

Copper . 

Dysprosium 

Erbium . 

Europium 

Fluorine 

Gadolinium 

Gallium 

Germanium 

Glucinum 

Gold . . 

Helium . 

Holmium 

Hydrogen 

Indium . 

Iodine . 

Iridium 

Iron . . 

Krypton 

Lanthanum 

Lead . . 

Lithium 

Lutecium 

Magnesium 

Manganese 

Mercury . 



Al 


27.10 


Molybdenum . 


. Mo 


Sb 


120.20 


Neodymium 


. Nd 


A 


39.90 


Neon .... 


. Ne 


As 


74.96 


Nickel .... 


. . Ni 


Ba 


137.37 


Niton .... 


. Nt 


Bi 


208.00 


Nitrogen . . . 


. N 


B 


10.90 


Osmium . . . 


. Os 


Br 


79.92 


Oxygen . . . 


. . O 


Cd 


112.40 


Palladium . . 


. Pd 


Cs 


132.81 


Phosphorus . . 


. . P 


Ca 


40.07 


Platinum . . 


. Pt 


C 


12.005 


Potassium . . 


. K 


Ce 


140.25 


Praseodymium 


. Pr 


CI 


35.46 


Radium . . . 


. Ra 


Cr 


52.00 


Khodium . . . 


. Rh 


Co 


58.97 


Rubidium . . 


. Rb 


Cb 


93.10 


Ruthenium . . 


. Ru 


Cu 


63.57 


Samarium . . 


. Sa 


By 


162.50 


Scandium . . 


. Sc 


Er 


167,70 


Selenium . . 


. Se 


Eu 


152.00 


Silicon . . . 


. Si 


F 


19.00 


Silver .... 


• Ag 


Gd 


157.30 


Sodium . . . 


. Na 


Ga 


70.10 


Strontium . . 


. Sr 


Ge 


72.50 


Sulfur .... 


. S 


Gl 


9.10 


Tantalum . . 


. . Ta 


Au 


197.20 


Tellurium . . 


. Te 


He 


4.00 


Terbium . . . 


. Tb 


Ho 


163.50 


Thallium . . . 


. Tl 


H 


1.008 


Thorium . . . 


. . Th 


In 


114.80 


Thulium . . . 


. . Tm 


I 


126.92 


Tin ..... 


. Sn 


Ir 


193.10 


Titanium . . 


. Ti 


Fe 


55.84 


Tungsten . . 


. W 


Kr 


82.92 


Uranium . . . 


. U 


La 


139.00 


Vanadium . . 


. . V 


Pb 


207.20 


Xenon .... 


. Xe 


Li 


6.94 


Ytterbium . . 


. Yb 


Lu 


175.00 


Yttrium . . . 


. Yt 


Mg 


24.32 


Zinc .... 


. Zn 


Mn 


54.93 


Zirconium . . 


. Zr 


Hg 


200.60 







WEIGHT IN GRAMS OF 1 LITER OF VARIOUS GASES UNDER STAND- 
ARD CONDITIONS AND BOILING POINTS UNDER PRESSURE OF 
760 MILLIMETERS 

Weight Boiling 
Name of 1 Liter Point 

Hydrogen chloride 1.6398 -83.1° 

Hydrogen fluoride 0.893 +19.4° 

Hydrogen sulfide 1.5392 -61.6° 

Methane .... 0.7168 -164.0° 

Nitric oxide . . 1.3402 -153.0° 

Nitrogen .... 1.2507 -195.7° 

Nitrous oxide . . 1.9777 -89.8° 

Oxygen .... 1.4290 -183.0° 

Sulfur dioxide . 2.9266 -8.0° 





Weight 


Boiling 


Name 


of 1 Liter 


Point 


Acetylene . . . 


1.1621 


-83.8° 


Air 


1.2928 




Ammonia . . . 


0.7708 


-33.5° 


Argon 


1.7809 


- 186.0° 


Carbon dioxide . 


1.9768 


-78.2° 


Carbon monoxide 


1.2504 


-190.0° 


Chlorine . . . . 


3.1674 


-33.6° 


Helium . . . . 


0.1782 


-268.7° 


Hydrogen . . . 


0.08987 


-252.7° 



DISPLACEMENT (ELECTROCHEMICAL) SERIES 



1. Potassium 

2. Sodium 

3. Lithium 

4. Calcium 

5. Magnesium 

6. Aluminium 



7. Manganese 

8. Zinc 

9. Chromium 

10. Iron 

11. Cobalt 

12. Nickel 



13. Tin 

14. Lead 

15. Hydrogen 
T6T Copper 

17. Arsenic 

18. Bismuth 



19. Antimony 

20. Mercury 

21. Silver 

22. Platinum 

23. Gold 



RELATION BETWEEN ENGLISH AND METRIC CONSTANTS 



1 pound (troy) 

1 ounce (troy) 

1 pound (avoirdupois) 

1 ounce (avoirdupois) 

1 kilogram 

1 kilogram 

1 liter 

1 gallon 

1 cubic centimeter 

1 cubic inch 

1 cubic foot 

1 centimeter 

1 meter 



= 373.24 grams 

= 31.10348 grams 

= 453.59 grams 

= 28.3495 grams 

= 2.67923 pounds (troy) 

= 2.20462 pounds (avoirdupois) 

= 1.05668 United States quarts 

= 3. 78543 liters 

= 0.0610 cubic inch 

= 16.3872 cubic centimeters 

= 28,320 cubic centimeters 

= 0.3937 inch 

= 39.37 inches 



LIBRARY OF CONGRESS 




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